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

Acid-Base Titrations & Indicators

In 2009, Australia's Therapeutic Goods Administration recalled a batch of Mylanta antacid tablets after in-house titration showed active base content was 18% below the labelled 680 mg per tablet — confirming that acid-base titration remains the legal standard for verifying what a medicine actually contains.

Today's hook: In 2009, the Australian TGA found that a recalled Mylanta antacid batch contained only 556 mg of active base per tablet — 18% less than the 680 mg labelled. The verification method was acid-base titration: a standard HCl solution was added from a burette until phenolphthalein changed colour, and stoichiometry did the rest. How does a single colour change in a flask translate into a legally defensible concentration value?

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

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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
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 for this lesson

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

Equivalence pointThe point in a titration where stoichiometrically equivalent amounts of acid and base have reacted.

EndpointThe point at which the indicator changes colour; ideally coincides with the equivalence point.

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

Primary standardA highly pure, stable substance used to prepare or check the concentration of a standard solution (e.g., Na₂CO₃, oxalic acid).

Indicator selectionChoose an indicator whose pH transition range overlaps the steep portion of the titration curve near the equivalence point.

Cross-lesson links: Titration calculations here use the mole-ratio method from Module 5. The equivalence-point concept reappears in L12 (drug pKa and the Henderson-Hasselbalch equation). Back titration of antacid tablets connects forward to L11–L12 (aspirin and paracetamol pharmacology).

Core Content

What Titration Measures

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.

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Essential Formulas

n = cV Moles = concentration × volume in litres

c = n / V Concentration = moles divided by volume in litres

At equivalence: mole ratio follows the balanced equation For 1:1 reactions, moles acid reacted = moles base reacted

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.

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

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?

Calculating Unknown Concentration

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 — Finding Concentration of Unknown Base

Given: 24.60 mL of 0.1000 mol L^-1 HCl(aq) neutralises 25.00 mL of NaOH(aq).

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

Find: Concentration of NaOH(aq).

Method: First calculate moles of HCl.

n(HCl) = cV = 0.1000 × 0.02460 = 0.002460 mol

The equation ratio is 1:1, so:

n(NaOH) = 0.002460 mol

Now calculate concentration of NaOH in 25.00 mL = 0.02500 L.

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

Answer: The sodium hydroxide concentration is 0.0984 mol L^-1.

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?

Back Titration for Antacid Tablets

Add known excess, then titrate what is left over

We just saw how a standard titration calculates unknown concentration from a single colour-change endpoint. 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.

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 of an Antacid Tablet

Given: A crushed antacid tablet is treated with 50.00 mL of 0.2000 mol L^-1 HCl(aq). The excess acid requires 18.40 mL of 0.1000 mol L^-1 NaOH(aq) for neutralisation. Active ingredient: NaHCO₃(s).

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

Find: Moles and mass of NaHCO₃ in the tablet.

Method: Calculate total moles of HCl added.

n(initial HCl) = 0.2000 × 0.05000 = 0.01000 mol

Use the NaOH titre to find excess HCl remaining.

n(NaOH) = 0.1000 × 0.01840 = 0.001840 mol

Because HCl and NaOH react 1:1:

n(excess HCl) = 0.001840 mol

So the acid that reacted with the tablet was:

n(HCl reacted) = 0.01000 − 0.001840 = 0.008160 mol

NaHCO₃ reacts with HCl in a 1:1 ratio, so:

n(NaHCO₃) = 0.008160 mol

Molar mass of NaHCO₃ = 84.01 g mol^-1.

m = nM = 0.008160 × 84.01 = 0.685 g

Answer: The tablet contains 0.008160 mol of NaHCO₃, which is 0.685 g.

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

Indicators, Endpoint and Equivalence Point

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 = 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.

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 from a burette to a solution of HCl containing phenolphthalein. Just before the equivalence point the solution is colourless. Predict: what will you see the instant one extra drop of NaOH is added past the equivalence point, and why?

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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?

Check Your Understanding

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

1. C — equivalence point is stoichiometric; endpoint is the observed indicator colour change.

2. B — only concordant titres should be averaged; the rough titre is excluded.

3. D — correct calculation using moles then concentration.

4. A — weak acid-strong base requires an indicator that changes in the basic range.

5. B — the second titration measures the excess reagent left over.

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 know the 2009 TGA Mylanta recall story — how has your understanding changed?

How would the TGA chemist have used a standard HCl solution and back titration to prove the tablets contained only 556 mg active base instead of the labelled 680 mg?
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.

I've completed this lesson

Review

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.