Year 12 Chemistry Module 8 ⏱ ~35 min 5 MC · 3 Short Answer Lesson 2 of 16 Core

Gravimetric Analysis

Turn an invisible dissolved ion into a pure solid, weigh it, and use stoichiometry to defend what was present in the original sample.

Today's hook: A wastewater sample looks completely clear, yet after barium ions are added a white solid appears. How can the mass of that solid reveal the sulfate concentration of the original water?
0/5TASKS
1

Connect: make invisible ions measurable

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

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
Case Entry Before Method

An environmental chemist adds barium ions to a wastewater sample and a fine white solid appears. After filtering and drying that solid, the chemist uses only its mass to judge the sulfate content of the original water.

  • Why might weighing a solid precipitate be enough to determine how much sulfate ion was in the sample?
  • What parts of the process could make the final measured mass too high or too low?
By the end of this lesson you will:

Know

  • The sequence dissolve → precipitate → digest → filter → wash → dry → cool → weigh
  • How gravimetric analysis determines the amount of analyte from precipitate mass
  • Which precipitating agents are suitable for common ions

Understand

  • Why an insoluble, pure precipitate is essential for reliable gravimetric analysis
  • How stoichiometry links precipitate mass to original analyte amount
  • How incomplete precipitation, co-precipitation, incomplete drying and filtration losses change results

Can Do

  • Calculate analyte mass or percentage composition from precipitate data
  • Select a suitable precipitating agent for Cl-, SO42- and CO32-
  • Interpret gravimetric results and judge whether a data set is reliable
Vocabulary flip cards+10 XP

Say the definition before you flip. Then choose “Got it” or “Again”.

0 / 6 mastered
Term 1Gravimetric analysisClick to reveal ↻
DefinitionDetermining analyte amount by converting it to a pure, weighable product of known formula.
Term 2Precipitating agentClick to reveal ↻
DefinitionA reagent that converts the target species into a low-solubility solid.
Term 3DigestionClick to reveal ↻
DefinitionAllowing a precipitate to stand, often warm, so larger and purer crystals form.
Term 4Constant massClick to reveal ↻
DefinitionSuccessive dry, cool and weigh cycles give the same mass within the required precision.
Term 5Co-precipitationClick to reveal ↻
DefinitionImpurities are carried down or trapped with the desired precipitate, increasing its mass.
Term 6DesiccatorClick to reveal ↻
DefinitionA dry sealed container used to cool a hot precipitate without moisture uptake.
Cross-lesson links: Gravimetric analysis relies on the same mole-conversion pathway used in titration (L01). Identifying which ions form insoluble precipitates connects directly to L03 (qualitative analysis). Environmental applications overlap with L06–L09 (water quality and contamination monitoring).
Safety: Wear safety glasses and use tongs when handling hot crucibles and drying ovens, as they can cause serious burns. Many precipitating reagents (for example silver nitrate and barium chloride) are corrosive or toxic, so avoid skin contact and wash your hands after use. Place silver and heavy-metal residues in labelled waste containers, not down the sink.
Core Content
2

Explain: mass becomes chemical evidence

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What Gravimetric Analysis Measures

Use the mass of a known precipitate to infer the mass of an unknown analyte

Gravimetric analysis is chemistry reduced to its most disciplined form: if you can isolate a pure precipitate of known composition, its mass becomes a direct clue to the original sample.

In gravimetric analysis, the ion or compound of interest is converted into an insoluble precipitate with a known chemical formula. Once that precipitate has been collected and dried, its mass can be used to calculate the amount of analyte originally present.

For sulfate analysis, adding Ba2+(aq) produces barium sulfate, BaSO4(s), a very insoluble white precipitate:

Gravimetric analysis determines analyte amount from the mass of an insoluble precipitate of known formula. Core formula sequence: n = m/M on precipitate → mole ratio from balanced equation → m = nM on analyte → percentage composition = (mass of analyte / mass of sample) × 100.

Pause, copy the highlighted formula sequence into your book.

→
Core Gravimetric Logic
Ba2+(aq) + SO42-(aq) → BaSO4(s) One mole of sulfate ion produces one mole of barium sulfate precipitate.
n = m / M Moles of precipitate = mass divided by molar mass
% composition = (mass of analyte / mass of sample) × 100 Once analyte mass is found, percentage composition follows directly.
HSC language: When describing gravimetric analysis, refer to conversion of the analyte into an insoluble precipitate of known composition, followed by isolation, drying and weighing so stoichiometric calculations can determine the original amount present.
Gravimetric analysis 5-step sequence migrated to sequence template

Gravimetric analysis succeeds only if the precipitate is fully formed, completely collected, thoroughly dried, and then weighed accurately. Each step contributes directly to the final chemical inference.

In gravimetric analysis of sulfate ions, which formula gives moles of the BaSO4 precipitate from its measured mass?
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Represent: the full practical sequence

2
The Practical Process

Eight controlled stages, one defensible mass

We just saw that gravimetric analysis links precipitate mass to analyte amount through mole ratios. That raises a question: how does the chemist produce a pure, collectable precipitate reliably? This card answers it with the full practical sequence.

A gravimetric result is only as good as the technique used to isolate the solid. The chemistry may be simple, but the method is unforgiving.

  1. Dissolve: Ensure the analyte is fully in solution so it can react completely.
  2. Precipitate: Add an appropriate reagent to form a low-solubility precipitate.
  3. Digest: Warm and allow the solid to stand so larger, purer crystals form.
  4. Filter: Separate the solid from the liquid without losing precipitate.
  5. Wash: Remove soluble ions adhering to the precipitate without dissolving it.
  6. Dry: Remove water so the mass measured is the mass of the precipitate, not liquid trapped with it.
  7. Cool: Cool in a desiccator so the solid does not absorb moisture before weighing.
  8. Weigh: Repeat heat, cool and weigh cycles until constant mass is reached.

If any step is incomplete, the final mass no longer represents the true amount of precipitate formed. Gravimetric analysis therefore depends on both chemical selectivity and careful laboratory technique.

Full gravimetric sequence: dissolve → precipitate → digest → filter → wash → dry → cool → weigh. Repeat drying, cooling and weighing until constant mass. Loss of solid lowers the result; retained water or contamination raises it.

Pause, copy the highlighted sequence into your book.

Industry anchor: In wastewater testing, gravimetric analysis is especially useful when the target ion forms a stable insoluble salt. For sulfate discharge monitoring, a well-dried BaSO4(s) precipitate can provide an inexpensive and defensible measure of contamination.
Why are digestion and washing included before the precipitate is dried?
4

Respond: choose a selective precipitate

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Choosing the Right Precipitating Agent

Solubility rules decide whether the method will work

We just saw the full practical sequence. That raises a question: which reagent do you add during precipitation, and why does it matter? This card answers it → the precipitating agent must produce an insoluble solid of known formula that is specific to the target ion.

A gravimetric method succeeds only when the reagent produces a precipitate that is both sufficiently insoluble and chemically specific.

Target ionPrecipitating reagentPrecipitate formedWhy it works
Cl-AgNO3(aq)AgCl(s)Silver chloride is insoluble and forms a distinct solid
SO42-BaCl2(aq) or Ba(NO3)2(aq)BaSO4(s)Barium sulfate has very low solubility
CO32-CaCl2(aq)CaCO3(s)Calcium carbonate precipitates from solution

A precipitating reagent must not simply "make a solid". It must form a precipitate with known composition, low solubility, and minimal side reactions with other ions in solution.

Key precipitating agents: Cl⁻ → AgNO₃(aq) → AgCl(s); SO₄²⁻ → BaCl₂(aq) → BaSO₄(s); CO₃²⁻ → CaCl₂(aq) → CaCO₃(s). Each precipitate must have known formula, low solubility, and high chemical specificity for the target ion.

Pause, copy the highlighted table into your book before the check.

Common error: "Any reagent that makes cloudiness is fine." Cloudiness alone is not enough. The precipitate must be identifiable, sufficiently insoluble, and suitable for accurate stoichiometric conversion to the analyte.
Which reagent is most suitable for precipitating sulfate ion in a gravimetric analysis?
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Apply: calculate back to the analyte

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Calculations from Precipitate Mass

Mass of precipitate → moles of precipitate → moles of analyte → mass or percentage

We just saw how choosing the right precipitating agent ensures a known compound forms. That raises a question: once you have a mass reading, what is the calculation path to the original analyte? This card answers it → convert precipitate mass to moles, apply the mole ratio, then calculate analyte mass and percentage.

The reliable way to solve gravimetric questions is to convert through moles. Do not jump straight from precipitate mass to percentage by intuition.

01 · balancemass of solidUse the dry precipitate at constant mass.
02 · convertn = m / MFind moles of the known precipitate.
03 · equationmole ratioConvert precipitate moles to analyte moles.
04 · reportm = nMFind analyte mass, then percentage if required.

For BaSO4(s), the mole ratio to SO42- is 1:1. That makes sulfate gravimetric analysis especially clean: one mole of precipitate corresponds to one mole of sulfate ion in the original sample.

Gravimetric calculation pathway: n(precipitate) = m/M → use mole ratio → m(analyte) = n × M(analyte) → % composition = (m(analyte) / m(sample)) × 100. For BaSO₄: M = 233.39 g mol⁻¹; for SO₄²⁻: M = 96.06 g mol⁻¹.

Pause, copy the highlighted calculation pathway into your book.

Worked example · sulfate by mass

Problem: A 0.500 g sample produces 0.350 g BaSO4. Find the mass percentage of sulfate. M(BaSO4) = 233.39 g mol-1; M(SO42-) = 96.06 g mol-1.

1
Find precipitate moles
n(BaSO4) = 0.350 / 233.39
= 0.00150 mol
The measured solid has a known formula and molar mass.
2
Apply the equation ratio
BaSO4 : SO42- = 1 : 1
n(SO42-) = 0.00150 mol
Each formula unit of precipitate contains one sulfate ion.
3
Find sulfate mass
m = nM = 0.00150 × 96.06
= 0.144 g
Convert analyte moles, not precipitate moles, to analyte mass.
4
Find mass percentage
% = (0.144 / 0.500) × 100
= 28.8%
Compare sulfate mass with the original sample mass.
Step 1 of 4 · reason before answer
A 0.500 g sample produces 0.350 g of BaSO4(s). Using M(BaSO4)=233.39, M(SO42-)=96.06, what is the sulfate percentage?
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Feedback: predict the direction of error

Beyond the syllabus. The gravimetric procedure, stoichiometry, drying to constant mass and error interpretation are core. Digestion, desiccator detail and co-precipitation mechanisms are extension — exam questions assess the method and the calculation.
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Sources of Error in Gravimetric Analysis

Why the final mass can be too low or too high

We just saw how to calculate from precipitate mass to percentage composition. That raises a question: what can go wrong with the precipitate mass itself? This card answers it → four distinct error sources, each with a predictable direction: too low leads to underestimation, too high leads to overestimation.

Imagine a chemist filters and dries a sulfate precipitate, places it on the balance, and reads 0.4382 g. That number looks precise, but if impurities co-precipitated, or if the filter paper ash was not removed, the number is wrong in ways the balance cannot detect. A gravimetric result looks objective because it ends with a balance reading, but that reading can still be wrong for several chemical reasons.

Incomplete precipitation Analyte remains dissolved, so too little solid is collected.mass ↓ · result underestimated
Co-precipitation Other ions or impurities are trapped with the solid.mass ↑ · result overestimated
Incomplete drying Water remains in or on the precipitate when weighed.mass ↑ · result overestimated
Filtration or transfer loss Some precipitate passes through or remains on apparatus.mass ↓ · result underestimated

Errors causing underestimation (mass too low): incomplete precipitation, filtration loss. Errors causing overestimation (mass too high): co-precipitation of impurities, incomplete drying. Always state the direction: "The measured mass is too high/low, so the analyte amount is overestimated/underestimated."

Pause, copy the highlighted error table into your book.

Must know: In HSC responses, always link the error to its direction. Saying "incomplete drying is an error" is not enough; you should say it causes the measured mass to be too high, so the analyte amount is overestimated.
Sources of error in gravimetric analysis: too-low results from incomplete precipitation (not all analyte becomes solid) or filtration loss (precipitate lost); too-high results from co-precipitation (impurities add mass) or incomplete drying (water remains).

The key HSC move is to connect each procedural error to its direction: low measured mass causes underestimation, while extra mass from water or impurities causes overestimation.

A precipitate is not dried fully before weighing. What is the most likely effect on the result?
Interactive Tool, Gravimetric Analysis Open fullscreen ↗
True or false? In gravimetric analysis, a precipitate that is not dried to constant mass will give a result that is too high, because the retained water is weighed as if it were analyte.
🔀Sort the Steps+7 XP
Arrange these steps of a gravimetric analysis to determine chloride ion concentration in a water sample.
Filter the precipitate quantitatively
Cool it in a desiccator
Dissolve the measured sample so the chloride ions are in solution
Dry the precipitate thoroughly
Digest the precipitate so larger, purer crystals form
Weigh, then repeat dry-cool-weigh cycles to constant mass
Add excess silver nitrate to precipitate AgCl(s)
Wash away soluble ions without dissolving AgCl
✓

Practice: defend a gravimetric result

Complete the Learn phase to unlock Practice.

Activities
A1
Calculating Sulfate from Wastewater Data

A 0.500 g dried wastewater residue was dissolved and treated with excess BaCl2(aq). The precipitated BaSO4(s) was filtered, dried and weighed in three trials:

TrialMass dry sample / gMass BaSO4 / gObservation
10.5000.348White precipitate, dried to constant mass
20.5000.351White precipitate, dried to constant mass
30.5000.392Sample removed from oven early; still slightly damp

1. Which trial should be excluded, and what specific procedural issue makes it unreliable?

2. Calculate the average valid mass of BaSO4(s).

3. Calculate the sulfate percentage in the dried residue using the average valid mass.

A2
Choosing Reagents and Diagnosing Errors

1. A chemist wants to determine chloride concentration in river water by gravimetric analysis. Which reagent should be added, and what precipitate forms?

2. A precipitate is weighed before it is fully dry. Explain the effect on the measured mass and the final analyte calculation.

3. During filtration, some of the solid passes through torn filter paper. Explain the effect on the result.

4. Why is Ba2+(aq) preferred over Na+(aq) for sulfate gravimetric analysis?

Check Your Understanding
5Q
Practice questions · Randomised bank

Answer five questions from the L02 bank. The options reshuffle, so reason from the method and chemistry.

SA
Short Answer Practice

1. Describe how a chemist would determine the percentage composition of sulfate in a wastewater sample using gravimetric analysis. In your answer, refer to precipitation, isolation of the solid, and calculation steps. 4 marks

2. Explain how incomplete precipitation and loss of precipitate on filtration would each affect the final calculated analyte content. 4 marks

3. Evaluate the suitability of gravimetric analysis for monitoring sulfate concentration in industrial wastewater. In your answer, refer to one strength of the method, one limitation or error risk, and whether the method provides enough evidence for environmental decision-making. 5 marks

Show All Answers

Activity 1

1. Exclude Trial 3 because the note states the precipitate was still slightly damp. Incomplete drying makes the measured mass too high, so sulfate content would be overestimated.

2. Average valid mass = (0.348 + 0.351) / 2 = 0.3495 g.

3. n(BaSO4) = 0.3495 / 233.39 = 0.00150 mol. Therefore n(SO42-) = 0.00150 mol. m(SO42-) = 0.00150 × 96.06 = 0.144 g. % sulfate = (0.144 / 0.500) × 100 = 28.8%.

Activity 2

1. Use AgNO3(aq). It forms AgCl(s), an insoluble silver chloride precipitate.

2. If the precipitate is not fully dry, extra water is included in the balance reading. The measured mass is too high, so the analyte amount is overestimated.

3. Losing precipitate during filtration makes the final mass too low, so the analyte amount is underestimated.

4. Ba2+(aq) forms insoluble BaSO4(s). Sodium sulfate remains soluble, so Na+(aq) would not produce a useful gravimetric precipitate.

Multiple Choice Explanations

Method sequence. The full controlled sequence is dissolve, precipitate, digest, filter, wash, dry, cool and weigh.

Sulfate reagent. BaCl2(aq) forms insoluble BaSO4(s).

Incomplete drying. Residual water adds mass and causes overestimation.

Co-precipitation. Trapped impurities increase measured mass.

Known outlier. A damp precipitate has a specific procedural cause for an artificially high result.

Short Answer Model Answers

Q1 (4 marks): The wastewater sample is dissolved so the sulfate ions are in solution. A solution containing Ba2+(aq), such as BaCl2(aq), is added to form BaSO4(s), an insoluble white precipitate. The precipitate is then filtered, dried thoroughly, and weighed. Its mass is converted to moles using n = m/M, and because the mole ratio between BaSO4 and SO42- is 1:1, the moles and mass of sulfate in the original sample can be calculated. Percentage composition is then found using (mass of sulfate / mass of original sample) × 100.

Q2 (4 marks): Incomplete precipitation means some analyte remains dissolved instead of forming the solid precipitate. This makes the measured precipitate mass too low, so the analyte content is underestimated. Loss of precipitate on filtration also reduces the final mass because some solid is physically lost before weighing. This again causes the analyte content to be underestimated. Although the causes are different, both errors lower the measured precipitate mass and therefore the calculated result.

Q3 (5 marks): Gravimetric analysis is suitable for monitoring sulfate in industrial wastewater because sulfate forms a very insoluble precipitate, BaSO4(s), allowing a direct stoichiometric link between precipitate mass and sulfate amount. A major strength is that the method is simple, inexpensive and based on measurable mass rather than subjective colour intensity. However, it is vulnerable to procedural errors such as incomplete drying or co-precipitation, both of which can distort the result. Overall, gravimetric analysis provides strong evidence for environmental monitoring when the precipitate is pure and dried to constant mass, especially if repeat trials are consistent and any outliers are justified scientifically.

Return to Think First

Return to the wastewater sample from Think First. Now that you understand gravimetric analysis, explain how the white precipitate becomes defensible evidence for sulfate content.

  • How does the stoichiometry of BaSO4 precipitation let you calculate sulfate amount from a balance reading?
  • Which error sources would have made the gravimetric result too high or too low, and how would chemists design the procedure to minimise them?
  • Write one sentence explaining why BaSO4(s) is a suitable precipitate for sulfate analysis.
✓

Review: check what actually stuck

Quick Quiz
Review

State the eight stages of gravimetric analysis in order.

Which reagent precipitates SO42-, and why is it suitable?

A gravimetric result is too high. Name two possible causes.

Describe the calculation pathway from precipitate mass to percentage composition.

Why does incomplete precipitation cause underestimation of the analyte?