Difficulty: Introductory to Intermediate | Prerequisites: General chemistry foundations (stoichiometry, molarity, solubility rules, Ksp). Familiarity with lab balances and basic glassware is helpful.
Gravimetric analysis is one of the oldest and most straightforward techniques in quantitative analytical chemistry. It works by converting the analyte (the thing you want to measure) into a weighable precipitate with known composition. In this experiment, chloride ions in an unknown salt are precipitated as silver chloride (AgCl) using excess silver nitrate, and the mass of that precipitate is used to back-calculate how much chloride was in the original sample. If you are comfortable with stoichiometry and the idea of a limiting reagent, this procedure is a natural next step into real quantitative lab work.
Dissolve an unknown chloride salt in water, add excess AgNO₃ to precipitate all chloride as AgCl, filter and dry the precipitate to constant weight, then use the mass of AgCl and stoichiometry to calculate the percent chloride in your original sample. The method relies entirely on accurate mass measurements, so drying, complete collection of precipitate, and propagation of weighing uncertainties are the make-or-break details.
Gravimetric analysis
A quantitative analytical method in which the amount of an analyte is determined by measuring the mass of a substance of known composition that has been chemically separated from the sample.
In simple terms, this means you figure out how much of something is in your sample by weighing a solid product you made from it.
Analyte
The substance or chemical component being identified or measured in an analysis.
Think of it as the thing you are trying to find or quantify, in this case, chloride ions.
Precipitate
An insoluble solid that forms when two solutions are mixed or when a solution is supersaturated with a compound.
In simple terms, it is the solid that crashes out of solution during a reaction. Here, AgCl is the precipitate.
Double displacement reaction (double replacement, metathesis)
A reaction in which the cations and anions of two compounds swap partners to form two new compounds.
Think of it as two pairs of dance partners switching: XCl + AgNO₃ becomes XNO₃ + AgCl.
Solubility product constant (Ksp)
The equilibrium constant for the dissolution of a sparingly soluble ionic solid in water, equal to the product of the ion concentrations each raised to the power of their stoichiometric coefficients.
In simple terms, Ksp tells you how much of an "insoluble" salt will still dissolve. A very small Ksp (AgCl: 1.8 × 10⁻¹⁰) means almost none of it dissolves, which is exactly what makes the method work.
Constant weight (drying to constant weight)
The point at which successive weighings of a dried sample agree within the balance's precision, indicating all moisture has been removed.
Think of it as heating and weighing, then heating and weighing again, until the number stops changing.
Weighing by difference
A technique where the mass of a substance is found by subtracting the mass of the empty container from the mass of the container plus substance.
In simple terms: (crucible + AgCl) minus (empty crucible) = mass of AgCl.
Propagation of uncertainty (error propagation)
A mathematical method for estimating the combined uncertainty in a calculated result based on the individual uncertainties of each measured quantity used in the calculation.
Think of it as working out how much your final answer could be off, given that every measurement you fed into it was off by a small amount.
The unknown salt (XCl) is dissolved in deionised water and reacted with excess 0.15 M silver nitrate (AgNO₃).
The reaction is a double displacement: XCl(s) + AgNO₃(aq) → XNO₃(aq) + AgCl(s)
AgCl is the insoluble precipitate (Ksp = 1.8 × 10⁻¹⁰), which forms as a white solid.
0.1 M nitric acid (HNO₃) is added to the solution. It acts as a catalyst and is not part of the stoichiometry, so it does not appear in the balanced equation.
AgNO₃ is deliberately added in excess so that all chloride ions are consumed. Leftover AgNO₃ stays in solution and does not interfere.
Drying the unknown salt to constant weight: Removes all moisture so the starting mass reflects only the salt, not trapped water.
Using deionised water to dissolve the salt: Prevents ions from tap water interfering with the precipitation reaction.
Using excess AgNO₃: Ensures every Cl⁻ ion reacts. The excess stays dissolved and washes away.
Washing precipitate with dilute 0.01 M HNO₃ (not pure water): Maintains a high electrolyte concentration, which keeps the AgCl particles clumped (coagulated) rather than dispersed as a colloid. Clumped particles are far easier to collect completely.
Scraping with a rubber policeman: Recovers precipitate stuck to the beaker walls. Leaving any behind directly lowers the calculated percent chloride.
Drying precipitate at 110 °C to constant weight: Drives off all moisture so the final mass is pure AgCl.
The molar ratio of AgCl to Cl⁻ in the reaction is 1:1.
From the mass of AgCl collected, calculate moles of AgCl: mol AgCl = mass AgCl / molar mass AgCl (143.321 g/mol).
Because the ratio is 1:1, mol Cl⁻ = mol AgCl.
Convert to mass of chloride: mass Cl = mol Cl⁻ × atomic mass of Cl (35.45 g/mol).
Percent chloride: %Cl = (mass Cl / mass of unknown sample) × 100.
Three crucibles gave percent chloride values of 66.75%, 71.87%, and 75.77%.
Average: 71.46%.
Variation from the average was within roughly ±6%, attributed mainly to incomplete collection of precipitate.
Balanced reaction:
XCl(s) + AgNO₃(aq) → XNO₃(aq) + AgCl(s)
Moles of AgCl:
mol AgCl = mass of AgCl (g) / 143.321 (g/mol)
Mass of chloride:
mass Cl (g) = mol AgCl × 35.45 (g/mol)
(The 1:1 stoichiometry means mol Cl⁻ = mol AgCl, so you go straight from moles of precipitate to mass of chloride.)
Percent chloride:
%Cl = (mass Cl / mass of unknown sample) × 100
Ksp expression for AgCl:
Ksp = [Ag⁺][Cl⁻] = 1.8 × 10⁻¹⁰
Gravimetric chloride determination is used in water-quality testing to measure chloride levels in drinking water and wastewater, since elevated chloride can corrode pipes and harm aquatic life. The same principle applies in cement and concrete testing, where chloride content above certain thresholds accelerates reinforcement corrosion. Pharmaceutical labs use gravimetric methods to verify the purity and composition of chloride-containing drug salts.
Students often think HNO₃ is a reactant in the precipitation. It is not. It acts as a catalyst and keeps the solution acidic to prevent hydroxide co-precipitation. It does not appear in the balanced equation.
Students sometimes assume that using excess AgNO₃ will cause errors. It will not, because the excess stays dissolved and is washed away. The point of the excess is to ensure every Cl⁻ ion is captured.
Washing the precipitate with pure water instead of dilute HNO₃ seems logical, but pure water causes peptisation: the AgCl particles disperse into a colloid and pass through the filter, leading to loss of product.
A common assumption is that one weighing after drying is enough. Drying to constant weight means repeating the heat-and-weigh cycle until successive masses agree, confirming all moisture is gone.
⚠️ Be able to write the balanced double displacement equation for any metal chloride reacting with AgNO₃.
⚠️ Know the Ksp of AgCl (1.8 × 10⁻¹⁰) and what a small Ksp means for solubility and precipitate completeness.
⚠️ Expect a calculation question: given mass of unknown and mass of AgCl, compute percent chloride. The 1:1 mole ratio and the unit conversions are the core of it.
⚠️ Be prepared to explain why dilute acid (not water) is used to wash the precipitate.
⚠️ Understand why the sample and the precipitate are both dried to constant weight, and what would go wrong if they were not.
True or False: HNO₃ appears in the balanced equation for this precipitation reaction. (False. It is a catalyst, not a reactant.)
True or False: Washing the precipitate with pure water is preferable to washing with dilute HNO₃. (False. Pure water causes peptisation and loss of precipitate.)
Fill in the blank: The mole ratio of AgCl to Cl⁻ in this reaction is ___. (1:1)
True or False: A larger Ksp value would make this method more reliable. (False. A smaller Ksp means more complete precipitation, which is what you want.)
Fill in the blank: "Drying to constant weight" means repeating the heat-and-weigh cycle until ___ . (successive masses agree within the balance's precision.)
Q: A student collects 0.4215 g of AgCl from 0.2000 g of an unknown chloride salt. What is the percent chloride in the sample?
A: mol AgCl = 0.4215 / 143.321 = 0.002941 mol. Mass Cl = 0.002941 × 35.45 = 0.1043 g. %Cl = (0.1043 / 0.2000) × 100 = 52.13%.
Q: Why is AgNO₃ added in excess rather than in a stoichiometric amount?
A: Using excess ensures that every Cl⁻ ion in solution reacts and precipitates as AgCl. If AgNO₃ were limiting, some chloride would remain dissolved and the calculated percent chloride would be too low.
Q: What role does 0.1 M HNO₃ play in the reaction mixture?
A: It keeps the solution acidic, which prevents the co-precipitation of silver hydroxide (AgOH) and ensures AgCl forms cleanly. It is not consumed in the reaction and does not appear in the balanced equation.
Q: A student washes the AgCl precipitate with pure deionised water and obtains a lower percent chloride than expected. Explain what likely happened.
A: Washing with pure water causes peptisation: the ionic strength of the wash drops, the AgCl particles disperse into a colloid, and some pass through the filter paper. The lost precipitate means less AgCl mass is recorded, giving a lower percent chloride.
Q: Explain why both the unknown salt and the final precipitate must be dried to constant weight.
A: Any residual moisture in the unknown inflates its apparent mass, making the calculated percent chloride too low. Residual moisture in the precipitate inflates its apparent mass, making the calculated percent chloride too high. Drying to constant weight removes this source of error in both directions.
This connects to solubility equilibria (Ksp calculations) because the success of the method depends on AgCl being nearly insoluble. If you are studying common-ion effects or Le Chatelier's principle, consider how the excess Ag⁺ shifts the equilibrium to drive precipitation further to completion.
Stoichiometry is the backbone of the calculation chain, so this lab reinforces mole-ratio reasoning from general chemistry. The error-propagation component ties into statistics and data-analysis topics you will encounter in physical chemistry or instrumental analysis courses.
Gravimetric analysis, gravimetric determination of chloride, AgCl precipitation, silver chloride precipitate, double displacement reaction, double replacement reaction, metathesis reaction, Ksp silver chloride, solubility product constant, weighing by difference, constant weight, drying to constant weight, propagation of uncertainty, error propagation, percent chloride calculation, quantitative analytical chemistry, rubber policeman, peptisation, precipitation reaction, analytical balance uncertainty