Gravimetric Analysis
In 2008, Chinese dairy company Sanlu was found to have adulterated infant formula with melamine — 2,563 tonnes were affected across the supply chain. Had inspectors applied gravimetric sulfate analysis to detect the urea-formaldehyde byproducts, the fraud's nitrogen-inflating trick would have been exposed months earlier.
Practise this lesson
Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
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?
Know
- The sequence dissolve → precipitate → filter → dry → 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
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.
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.
Dissolve → precipitate → filter → dry → weigh
We just saw that gravimetric analysis links precipitate mass to analyte amount through mole ratios. That raises a question: how does the chemist actually produce that precipitate reliably in the lab? This card answers it → the five-step practical sequence, where each step must be complete for the final mass to be valid.
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.
- Dissolve: Ensure the analyte is fully in solution so it can react completely.
- Precipitate: Add an appropriate reagent to form a low-solubility precipitate.
- Filter: Separate the solid from the liquid without losing precipitate.
- Dry: Remove water so the mass measured is the mass of the precipitate, not liquid trapped with it.
- Weigh: Measure the dry precipitate accurately and use stoichiometry to work backwards.
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.
Five steps of gravimetric analysis in order: Dissolve → Precipitate → Filter → Dry → Weigh. Drying must reach constant mass (re-weigh after further heating until the mass is unchanged). Any loss in filtration or any residual water changes the final mass and therefore the result.
Pause — copy the highlighted five steps into your book.
Solubility rules decide whether the method will work
We just saw the five-step practical process. That raises a question: which reagent do you add in step 2, 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 ion | Precipitating reagent | Precipitate formed | Why 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.
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.
- Measure the mass of dry precipitate formed.
- Calculate moles of precipitate using
n = m / M. - Use the balanced equation to find moles of analyte.
- Convert moles of analyte into mass using
m = nM. - If needed, calculate percentage composition from the original sample mass.
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.
Given: A 0.500 g sample produces 0.350 g of BaSO4(s).
Ba2+(aq) + SO42-(aq) → BaSO4(s)Molar mass of BaSO4 = 233.39 g mol-1. Molar mass of SO42- = 96.06 g mol-1.
Find: Mass and percentage of sulfate ion in the original sample.
Method: Calculate moles of BaSO4.
n(BaSO4) = m / M = 0.350 / 233.39 = 0.00150 molThe equation ratio is 1:1, so:
n(SO42-) = 0.00150 molConvert sulfate moles to sulfate mass.
m(SO42-) = nM = 0.00150 × 96.06 = 0.144 gNow find percentage composition of sulfate in the 0.500 g sample.
% sulfate = (0.144 / 0.500) × 100 = 28.8%Answer: The sample contains 0.144 g of sulfate ion, which is 28.8% by mass.
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.
| Error source | What happens | Effect on measured mass | Effect on result |
|---|---|---|---|
| Incomplete precipitation | Not all analyte forms the solid | Too low | Analyte amount underestimated |
| Co-precipitation | Other ions or impurities become trapped in the precipitate | Too high | Analyte amount overestimated |
| Incomplete drying | Water remains in or on the precipitate | Too high | Analyte amount overestimated |
| Loss on filtration | Some precipitate passes through or is left behind | Too low | Analyte amount 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.
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.
Complete the Learn phase to unlock Practice.
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:
| Trial | Mass dry sample / g | Mass BaSO4 / g | Observation |
|---|---|---|---|
| 1 | 0.500 | 0.348 | White precipitate, dried to constant mass |
| 2 | 0.500 | 0.351 | White precipitate, dried to constant mass |
| 3 | 0.500 | 0.392 | Sample 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.
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?
1. Which sequence best describes the core gravimetric analysis process?
2. Which reagent is most suitable for precipitating sulfate ion in a gravimetric analysis?
3. A precipitate is not dried fully before weighing. What is the most likely effect?
4. Which statement best describes co-precipitation?
5. A student calculates sulfate content from all three wastewater trials, including the damp precipitate result. Why is this poor analytical 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
1. B — the correct order is dissolve, precipitate, filter, dry, then weigh.
2. D — BaCl2(aq) forms insoluble BaSO4(s).
3. A — incomplete drying adds water mass and makes the result too high.
4. C — co-precipitation means impurities are trapped with the precipitate, increasing measured mass.
5. B — the damp trial is a known outlier with a specific chemical reason for being too high.
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 the 2008 Sanlu melamine scandal. Now that you understand gravimetric analysis, how would a properly applied precipitate test have caught the fraud earlier?
- How does the stoichiometry of BaSO4 precipitation let you calculate the analyte amount from a single balance reading — and why would that calculation have exposed the inflated "protein" content?
- 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 the preferred precipitating agent for sulfate analysis.
State the five steps 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?