Stoichiometry (College Board AP® Chemistry): Exam Questions

6 H⁺(aq) + 2 MnO₄⁻(aq) + 5 H₂C₂O₄ (aq) → 10 CO₂ (g) + 8 H₂O(l) + 2 Mn²⁺(aq)

A student dissolved a 0.139 g sample of oxalic acid, H₂C₂O₄ , in water in an Erlenmeyer flask. Then the student titrated the H₂C₂O₄ solution in the flask with a solution of KMnO₄ , which has a dark purple color. The balanced chemical equation for the reaction that occurred during the titration is shown above.

Identify the species that was reduced in the titration reaction. Justify your answer in terms of oxidation numbers.

The student used a 50.0 mL buret to add the KMnO₄(aq) to the H₂C₂O₄(aq) until a faint lavender color was observed in the flask, an indication that the end point of the titration had been reached. The initial and final volume readings of the solution in the buret are shown below. Write down the initial reading and the final reading and use them to determine the volume of KMnO₄(aq) that was added during the titration.

Given that the concentration of KMnO₄(aq) was 0.0235 M, calculate the number of moles of MnO₄⁻ ions that completely reacted with the H₂C₂O₄ .

The student proposes to perform another titration using a 0.139 g sample of H₂C₂O₄ , but this time using 0.00143 M KMnO₄(aq) in the buret. Would this titrant concentration be a reasonable choice to use if the student followed the same procedure and used the same equipment as before? Justify your response.

A chemist reacts barium chloride solution with sodium sulfate solution to produce solid barium sulfate and aqueous sodium chloride, according to the balanced equation below:

BaCl₂ (aq) + Na₂SO₄ (aq) ⟶ BaSO₄ (s) + 2NaCl (aq)

A solution is prepared by mixing 75.0 mL of 0.250 M barium chloride with 60.0 mL of 0.350 M sodium sulfate.

Determine the limiting reactant.

Calculate the maximum mass of barium sulfate that can form.

A student filters the barium sulfate and dries it, finding the final mass of the solid to be 6.05 g. Calculate the percent yield.

Explain how using excess water during rinsing of the solid could affect the percent yield measurement.

A student performs a titration to determine the concentration of Fe²⁺ (aq)) using 0.0200 M MnO₄^- (aq). The balanced equation for the reaction is:

5Fe²⁺ (aq) + MnO₄⁻ (aq) + 8H⁺ (aq) ⟶ 5Fe³⁺ (aq) + Mn²⁺ (aq) + 4H₂O (l)

In one trial, 25.00 mL of FeSO₄ (aq) is titrated with 18.75 mL of 0.0200 M MnO₄⁻ (aq).

i) Calculate the number of moles of MnO₄⁻ (aq) added during the titration.

ii) Calculate the number of moles of FeSO₄ (aq) that reacted.

Determine the concentration of the FeSO₄ (aq) solution in mol/L.

The student repeats the experiment and obtains a significantly lower concentration of FeSO₄ (aq). Suggest one procedural error that could account for this result.

A student analyzes an unknown hydrogencarbonate, XHCO₃ (s). A 1.69 g sample of XHCO₃ is dissolved in water to make 100.0 mL of solution. A 25.00 mL portion of this solution is titrated with 0.100 M HCl (aq). The balanced equation for the reaction is:

XHCO₃ (aq) + HCl (aq) ⟶ XCl (aq) + CO₂ (g) + H₂O (l)

In one trial, 15.60 mL of 0.100 M HCl (aq) is required to react completely with the 25.00 mL sample of the XHCO₃ solution.

i) Calculate the number of moles of HCl (aq) added during the titration.

ii) Calculate the number of moles of XHCO₃ (aq) in the original 100.0 mL solution.

i) Determine the molar mass of XHCO₃ (s) based on the 2.50 g sample.

ii) Identify the element X in XHCO₃ (s).

A student investigates the thermal decomposition of sodium bicarbonate in a sealed flask at 398 K.

2NaHCO₃ (s) → Na₂CO₃ (s) + H₂O (g) + CO₂ (g)

The student heats 8.40 g of NaHCO₃ in an evacuated 2.00 L container. After the reaction reaches completion, the total pressure of the gaseous products is measured to be 1.22 atm at 398 K.

Calculate the number of moles of NaHCO₃ placed in the container.

Determine the theoretical number of moles of gas produced from complete decomposition.

Use the ideal gas law to calculate the total number of moles of gas produced in the reaction.

Compare your answers to parts (b) and (c). Calculate the percent yield of the decomposition reaction.

Suggest one possible reason why the percent yield is less than 100%, even though all of the NaHCO₃ was heated.

After the decomposition of Na₂CO₃, the solid residue remaining in the container is dissolved in water and reacted with excess HCl (aq).

Write a balanced molecular equation for the reaction that occurs between Na₂CO₃ and HCl.

The student bubbles the gas produced from part (f) through limewater, (Ca(OH)₂ (aq)).

Describe the expected observation and identify the gas responsible for the change.

Instead of collecting the gas, the student titrates the resulting solution with silver nitrate (AgNO₃).

Write the balanced net ionic equation for the reaction, and
explain how this titration could be used to confirm the amount of Na₂CO₃ present in the original sample.

A student is given a 25.0 mL sample of a solution of an unknown monoprotic acid and asked to determine the concentration of the acid by titration. The student uses a standardized solution of 0.110 M NaOH(aq), a buret, a flask, an appropriate indicator, and other laboratory equipment necessary for the titration.

The images below show the buret before the titration begins (below left) and at the end point (below right). What should the student record as the volume of NaOH(aq) delivered to the flask?

Based on the given information and your answer to part (a), determine the value of the concentration of the acid that should be recorded in the student's lab report.

In a second trial, the student accidentally added more NaOH(aq) to the flask than was needed to reach the end point, and then recorded the final volume. Would this error increase, decrease, or have no effect on the calculated acid concentration for the second trial? Justify your answer.

A student investigates the reaction between hydrochloric acid and calcium carbonate to produce carbon dioxide gas, water, and calcium chloride, as shown in the balanced equation below:

CaCO₃ (s) + 2HCl (aq) ⟶ CaCl₂ (aq) + CO₂ (g) + H₂O (l)

The student places a 4.15 g sample of calcium carbonate into excess hydrochloric acid and collects the carbon dioxide gas produced in a syringe at 298 K and 1.50 atm.

The student records the following data:

Calculate the maximum theoretical volume of carbon dioxide gas that can be produced from 3.80 g of CaCO₃ at 25 ^oC and 101.3 kPa.

Using the axes provided below, plot a graph of volume of CO₂ (y-axis) against time (x-axis) using the data in the table. Label any key features of the graph.

Based on the graph and data, justify whether the reaction went to completion and how the rate of reaction changed over time. Support your answer with evidence.

Sodium azide (NaN₃) is produced for use in car airbags by reacting sodium amide (NaNH₂) with dinitrogen monoxide gas (N₂O) according to the balanced equation below:

2NaNH₂ (s) + N₂O (g) ⟶ NaN₃ (s) + NaOH (aq) + NH₃ (g)

A factory needs 600 g of sodium azide. The reaction has a 93.0% yield.

Calculate the theoretical mass of sodium amide required to produce 600.0 g of sodium azide.

Determine the actual mass of sodium amide that must be used.

A student analyzes an impure sample of an unknown carbonate compound with the formula X₂CO₃ (s). The sample has a mass of 2.90 g and is 85.0% X₂CO₃ by mass. The sample is dissolved in water to make 100.0 mL of solution. A 25.00 mL portion of this solution is titrated with 0.200 M HCl (aq). The balanced equation for the reaction is:

X₂CO₃ (aq) + 2HCl (aq) ⟶ 2XCl (aq) + CO₂ (g) + H₂O (l)

In one trial, 18.40 mL of 0.200 M HCl (aq) is required to react completely with the 25.00 mL sample of the solution.

i) Calculate the number of moles of HCl (aq) added during the titration.

ii) Calculate the number of moles of X₂CO₃ (aq) in the original solution.

Identify the element X in X₂CO₃ (s). Justify your answer.

A 0.2186 g sample of an unknown metal carbonate with the formula M₂CO₃ (s) containing 92.0% M₂CO₃ by mass is dissolved in water to make 250.0 mL of solution. A 25.00 mL portion of this solution is titrated with 0.0150 M HCl (aq). The balanced equation for the reaction is:

M₂CO₃ (aq) + 2HCl (aq) ⟶ 2MCl (aq) + CO₂ (g) + H₂O (l)

25.30 mL of 0.0150 M HCl (aq) is required to react completely with the 25.00 mL sample of the solution.

Calculate the number of moles of M₂CO₃ (aq) in the original 250.0 mL solution.

Identify the element M in M₂CO₃ (s). Justify your answer.

An excess volume of HCl (aq) was added past the endpoint during the titration. Explain how this error would affect the calculated molar mass of M₂CO₃.