What you'll be able to do: Analyse gravimetric and redox titration data, including percent composition and sources of experimental error.
Two classical techniques turn the ideas in this unit into numbers on a lab bench: weighing a precipitate, and titrating with permanganate. Both are only as good as the technique, so understanding how errors propagate matters as much as the arithmetic.
These are recommended, not required. You can start this lesson at any time.
If an ion can be precipitated completely as a compound of known formula, then weighing that dry solid measures how much of the ion was present. The method needs no calibration curve and no standard solution, only an accurate balance and patient technique.
Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
Add a slight excess of precipitating agent to drive the reaction to completion, digest the mixture so crystals grow larger and filter more cleanly, wash the solid to remove adsorbed spectator ions, then dry to constant mass. Drying to constant mass, weighing repeatedly until the reading stops falling, is the step that proves the water is gone.
Convert the precipitate mass to moles, use the formula to find moles of the ion of interest, convert that to a mass, then divide by the original sample mass. The gravimetric factor is simply this chain expressed as a single multiplier.
percent = (mass of analyte / mass of sample) × 100
In acidic solution, MnO₄⁻ is reduced to the almost colourless Mn²⁺ while the analyte is oxidised. Because the purple colour disappears as it reacts, the first drop of excess titrant leaves a faint persistent pink, so no separate indicator is required. Sulfuric acid is used to acidify, because hydrochloric acid would itself be oxidised.
MnO₄⁻ + 8 H⁺ + 5 e⁻ → Mn²⁺ + 4 H₂O
For any mistake, ask whether the measured quantity ends up too high or too low, then follow it through the calculation. A precipitate that is not fully dried weighs too much, so the reported percentage is too high. A burette rinsed with water instead of titrant delivers dilute titrant, so more volume is needed and the calculated analyte appears too high.
Quantitative analysis based on measuring the mass of a pure isolated product.
Repeated drying and weighing until successive masses agree, showing all solvent has gone.
A titrant whose own colour change marks the endpoint, as permanganate does.
A slight excess of the precipitating agent ensures the analyte is removed completely.
Trace each procedural fault to whether the measured value becomes too high or too low.
percent = (mass of analyte / mass of sample) × 100
A 1.000 g sample containing sulfate is dissolved and treated with excess BaCl₂, yielding 0.750 g of dry BaSO₄. What is the percent sulfate by mass?
Try it first: Decide what the weighed solid tells you about moles before converting anything.
0 of 4 steps revealed.
A 25.00 mL sample of Fe²⁺ solution requires 18.60 mL of 0.0200 M KMnO₄ in acid to reach the first persistent pink. What is the iron concentration?
Try it first: Recall how many electrons permanganate accepts, since that sets the ratio.
0 of 4 steps revealed.
In the gravimetric experiment above, the student removes the filter paper before the precipitate reaches constant mass. What effect does this have on the reported percent sulfate?
0 of 4 steps revealed.
Why it's wrong: Chloride is itself oxidised by permanganate, consuming extra titrant.
Check instead: Acidify with sulfuric acid, which is not oxidised.
Why it's wrong: The solid includes the counter-ion added during precipitation.
Check instead: Convert to moles, apply the formula ratio, then convert back to the analyte mass.
Why it's wrong: The direction depends on which measurement was affected and where it appears in the calculation.
Check instead: Trace the specific quantity through each step of the arithmetic.
Why it's wrong: Residual solvent adds mass that is indistinguishable from product.
Check instead: Dry and reweigh until two successive masses agree.
No practice questions are available for this topic yet. You can still practice the whole unit.
Gravimetric analysis determines an unknown amount by converting it quantitatively into a pure, dry, insoluble solid and weighing it. Redox titration determines it by measuring the volume of an oxidising agent consumed, with permanganate acting as its own indicator. In both, the calculation runs from the measured quantity through moles and the mole ratio to the analyte, and every procedural fault can be traced to whether it makes the measured quantity too high or too low.
This lesson is original Chem Help content. No external sources were adapted.