Elements & Mixtures (College Board AP® Chemistry): Exam Questions

A chemist analyzes a mixture of two metals, Fe and Cu, in which the iron-to-copper mass ratio is 3:2. The metals are reacted with excess nitric acid (HNO_3​), forming soluble nitrate salts (Fe(NO₃​)₃​ and Cu(NO₃​)₂​). The chemist then evaporates the solution, leaving behind the dry metal nitrates. The following data is recorded:

Suggest an experimental method to confirm whether the original sample contained a mixture of metals. Justify your answer.

Determine the total mass of metals (Fe and Cu) in the sample, and explain why this mass remains unchanged before and after the reaction.

Using the Fe:Cu mass ratio provided, calculate the mass of Fe in the sample.

Calculate the percent by mass of Fe in the original sample.

A student analyzes a 1.00 g sample of compound X, which is known to contain only carbon and hydrogen. Tests are performed to determine the mass of each element in the sample.

The student finds the sample contains 0.857 g of carbon and 0.143 g of hydrogen. Calculate the number of moles of each element present in the sample.

Determine the molecular formula of compound X. The molar mass of compound X is 56.1 g/mol.

Using your answer from part (b), calculate the number of carbon atoms in the original 1.20 g sample of compound X.

Use the balanced chemical equation for the complete combustion of one mole of compound X to calculate the number of CO₂ molecules produced during combustion of the 1.20 g sample.

Another student proposes that the molecular formula of compound X is C₂H₄.

Explain why this formula is inconsistent with the data in the question.

State the molecular formula of a compound with the same empirical formula as compound X but a molar mass of 70.1 g/mol.

Explain one limitation of using only elemental mass data when analyzing compounds that may contain oxygen.

A student analyzes a 1.68 g sample of a metal oxide containing only an unknown metal (M) and oxygen. The oxide is reduced by heating in a stream of hydrogen gas, which removes the oxygen as water vapor. After complete reduction, 0.89 g of pure metal remains.

The student also analyzes a separate sample of the same oxide and finds the following:

• 0.0100 mol of metal ions

• 0.0150 mol of oxide ions

Determine the empirical formula of the compound using the molar quantities provided.

Would the empirical formula change if the student used a larger or smaller sample size? Justify your answer.

Calculate the percent by mass of oxygen in the compound.

Use your answer from parts (a) and (c) to determine the molar mass of the metal and suggest its identity.

Explain why identifying the metal based on its molar mass provides evidence that the sample is pure.

The student repeats the experiment and recovers only 0.83 g of metal from a 1.76 g sample of the oxide. Explain what this result suggests about the purity of the sample. Justify your answer.

Explain how a particle-level model of the metal oxide differs from that of a mixture.

A laboratory receives a powdered sample containing iron (Fe) and silicon dioxide (SiO₂). The sample is believed to be a mixture of these two substances. To analyze the sample, the following procedures are performed:

• A magnet is used to separate iron from the mixture.

• The remaining material is heated strongly with hydrofluoric acid (HF), which reacts with silicon dioxide according to the equation:

SiO₂ (s) + 4HF (aq) → SiF₄ (g) + 2H₂O (l)

• The total mass of the original sample is 15.00 g. 6.25 g remains after separation with the magnet.

Explain why the original sample is classified as a mixture and not a pure substance.

Calculate the percent by mass of iron in the original sample and explain what this value indicates about the composition of the mixture.

Describe how physical properties are used to separate the components of the mixture.

Explain how the expected observations of the reaction of the remaining material with hydrofluoric acid provide evidence that the other substance could be silicon dioxide.

A laboratory receives a 10.00 g solid sample suspected to be either:

• Pure copper(II) sulfate pentahydrate (CuSO₄·5H₂O), or

• A mixture of copper(II) sulfate pentahydrate (CuSO₄·5H₂O) and an inert, non-volatile impurity.

To analyze the sample, the following procedure is performed:

1. The sample is heated strongly in a crucible to remove all water of hydration. Heating continues until the mass remains constant, with a final mass of 7.85 g.

2. After cooling, the residue is stirred into distilled water. The entire solid dissolves, producing a blue solution.

Explain why the sample is heated until a constant mass is achieved before analysis.

Calculate the percent by mass of water lost during heating.

The theoretical percent by mass of water in pure CuSO₄·5H₂O is 36.1%. Use this value and your calculation from part (b) to justify whether the sample is a mixture or pure copper(II) sulfate pentahydrate.

Explain how this experiment demonstrates that mixtures can have variable composition, based on the mass data.

A municipal water treatment facility is analyzing a 500.0 g sample of river water suspected of containing both suspended solids (such as sand and organic matter) and dissolved ionic compounds.

To investigate, the following procedure is carried out:

1. The sample is filtered, leaving 7.3 g of solid residue on the filter paper.

2. The filtered water is evaporated to dryness, leaving 3.8 g of solid residue.

3. A portion of the evaporated residue is tested with silver nitrate (AgNO₃) solution. A white precipitate forms.

Explain why filtration is an appropriate first step in separating the components of this river water sample.

Calculate the percent by mass of suspended solids in the original river water sample.

What does the result of the test with silver nitrate suggest about the composition of the dissolved solids? Justify your answer.

Explain why the procedures used in this experiment do not allow full identification of all substances present in the river water sample.