Calorimetry (College Board AP® Chemistry): Exam Questions

Aluminum metal can be recycled from scrap metal by melting the metal to evaporate impurities.

Calculate the amount of heat needed to purify 1.00 mole of Al originally at 298 K by melting it. The melting point of Al is 933 K. The molar heat capacity of Al is 24 J/(mol·K), and the heat of fusion of Al is 10.7 kJ/mol.

The equation for the overall process of extracting Al from Al₂O₃ is shown. Which requires less energy, recycling existing Al or extracting Al from Al₂O₃ ? Justify your answer with a calculation.

Al₂O₃(s) → 2 Al(s) + O₂(g)             ΔH° = 1675 kJ / mol_rxn

2 C₃H₇OH (l) + 9O₂ (g)  →  6CO₂ (g) + 8H₂O (g)

A student performs an experiment to determine the enthalpy of combustion of 2-propanol, C₃H₇OH (l) , which combusts in oxygen according to the equation above. The student heats a sample of water by burning some of the C₃H₇OH (l) that is in an alcohol burner, as represented below. The alcohol burner uses a wick to draw liquid up into the flame. The mass of C₃H₇OH (l) combusted is determined by weighing the alcohol burner before and after combustion.

Data from the experiment are given in the table below.

Calculate the magnitude of the heat energy, in kJ, absorbed by the water. (Assume that the energy released from the combustion is completely transferred to the water.)

Based on the experimental data, if one mole of C₃H₇OH (l) is combusted, how much heat, in kJ, is released ? Report your answer with the correct number of significant figures.

A second student performs the experiment using the same mass of water at the same initial temperature.

However, the student uses an alcohol burner containing  C₃H₇OH (l) that is contaminated  with water, which is miscible with C₃H₇OH (l) . The difference in mass of the alcohol burner before and after the combustion in this experiment is also 0.55 g. Would the final temperature of the water in the beaker heated by the alcohol burner in this experiment be greater than, less than, or equal to the final temperature of the water in the beaker in the first student's experiment? Justify your answer.

A student investigates the enthalpy of solution, ΔH_soln, for two alkali metal halides, LiCl and NaCl. In addition to the salts, the student has access to a calorimeter, a balance with a precision of ±0.1 g, and a thermometer with a precision of ±0.1°C.

To measure ΔH_soln  for LiCl, the student adds 100.0 g of water initially at 15.0°C to a calorimeter and adds 10.0 g of LiCl(s) , stirring to dissolve. After the LiCl dissolves completely, the maximum temperature reached by the solution is 35.6°C.

i) Calculate the magnitude of the heat absorbed by the solution during the dissolution process , assuming that the specific heat capacity of the solution is 4.18 J/(g·°C). Include units with your answer.

ii) Determine the value of ΔH_soln for LiCl in kJ /mol_rxn .

To explain why  ΔH_soln  for NaCl  is different than that for LiCl, the student investigates factors that affect ΔH_soln  and finds that ionic radius and lattice enthalpy (which can be defined as the ΔH associated with the separation of a solid crystal into gaseous ions) contribute to the process . The student consults references and collects the data shown in the table below.

Write the complete electron configuration for the Na⁺ ion in the ground state.

Using principles of atomic structure, explain why the Na⁺ ion is larger than the Li⁺ ion .

Which salt, LiCl or NaCl, has the greater lattice enthalpy? Justify your answer.

Below is a representation of a portion of a crystal of LiCl. Identify the ions in the representation by writing the appropriate formulas ( Li⁺ or Cl⁻ ) in the boxes below.

The lattice enthalpy of LiCl is positive, indicating that it takes energy to break the ions apart in LiCl. However, the dissolution of LiCl in water is an exothermic process. Identify all particle-particle interactions that contribute significantly to the dissolution process being exothermic. For each interaction , include the particles that interact and the specific type of intermolecular force between those particles.

A student performs a calorimetry experiment to measure the enthalpy change for the dissolution of potassium chloride (KCl) in water. The following data is recorded:

• Mass of water: 100.0 g

• Initial temperature of water: 25.0 ^oC

• Final temperature of the solution: 21.5 ^oC

• Specific heat capacity of water: 4.18 J / g^oC

• Mass of KCl dissolved: 5.00 g

• Molar mass of KCl : 74.55 g / mol

Calculate the heat transferred during the dissolution process. Indicate whether the process is exothermic or endothermic.

Determine the enthalpy change (ΔH) for the dissolution of 1.00 mol of KCl.

Explain why a polystyrene cup is often used as a calorimeter in such experiments.

A student performs an experiment to determine the value of the enthalpy change, ΔH°_rxn, for the oxidation-reduction reaction represented by the balanced equation below.

Na₂S₂O₃(aq) + 4 NaOCl(aq) + 2 NaOH(aq) → 2 Na₂SO₄(aq) + 4 NaCl(aq) + H₂O(l)

Determine the oxidation number of Cl in NaOCl .

Calculate the number of grams of Na₂S₂O₃ needed to prepare 100.00 mL of 0.500 M Na₂S₂O₃(aq).

In the experiment, the student uses the solutions shown in the table below.

 Using the balanced equation for the oxidation-reduction reaction and the information in the table above, determine which reactant is the limiting reactant. Justify your answer.

The solutions, all originally at 20.0°C, are combined in an insulated calorimeter. The temperature of the reaction mixture is monitored, as shown in the graph below.

According to the graph, what is the temperature change of the reaction mixture?

The mass of the reaction mixture inside the calorimeter is 15.21 g.

i) Calculate the magnitude of the heat energy, in joules, that is released during the reaction. Assume that the specific heat of the reaction mixture is 3.94 J/(g·°C) and that the heat absorbed by the calorimeter is negligible.

ii) Using the balanced equation for the oxidation-reduction reaction and your answer to part (c), calculate the value of the enthalpy change of the reaction, ΔH°_rxn, in kJ/mol_rxn. Include the appropriate algebraic sign with your answer.

The student repeats the experiment, but this time doubling the volume of each of the reactants, as shown in the table below.

The magnitude of the enthalpy change, ΔH°_rxn, in kJ/mol_rxn, calculated from the results of the second experiment is the same as the result calculated in part (e)(ii) . Explain this result.

Write the balanced net ionic equation for the given reaction.

4 Fe(s) + 3 O₂( g) → 2 Fe₂O₃(s)      ∆H ° = − 1650 kJ / mol_rxn

A student investigates a reaction used in hand warmers, represented above. The student mixes Fe(s) with a catalyst and sand in a small open container. The student measures the temperature of the mixture as the reaction proceeds. The data are given in the following table.

The mixture (Fe(s), catalyst, and sand) has a total mass of 15.0 g and a specific heat capacity of 0.72  J/(g·°C). Calculate the amount of heat absorbed by the mixture from 0 minutes to 4 minutes.

Calculate the mass of Fe(s), in grams, that reacted to generate the amount of heat calculated in part (a).

In a second experiment, the student uses twice the mass of iron as that calculated in part (b) but the same mass of sand as in the first experiment. Would the maximum temperature reached in the second experiment be greater than, less than, or equal to the maximum temperature in the first experiment? Justify your answer.

A student is investigating the heat transfer associated with dissolving potassium chloride (KCl in water. The student places 10.0 g of solid KCl into 100.0 mL of water in a polystyrene calorimeter. The student records the following data:

State whether the dissolution of KCl in water is an exothermic or endothermic process. Justify your answer based on principles of heat transfer.

Determine the enthalpy change for the dissolution of KCl, in kJ per mole of solute.

A student investigates the enthalpy change for the neutralization of hydrochloric acid (HCl) with sodium hydroxide (NaOH) using a calorimeter. The student mixes equal volumes of 1.00 M HCl and 1.00 M NaOH and measures the temperature change.

Determine the enthalpy change per mole of water formed in this reaction.

In this experiment, it was assumed that the specific heat capacity of the solution is the same as water. In reality, the actual specific heat capacity of the solution is slightly lower than that of pure water. Explain how this assumption affects the calculated value of ΔH.

A student investigates the enthalpy change for the combustion of magnesium (Mg) in oxygen. Since direct combustion is not performed in the lab, the student determines the enthalpy change indirectly by conducting two calorimetry experiments involving reactions of magnesium and magnesium oxide with hydrochloric acid.

The specific heat capacity of the solution is assumed to be 4.18 J/g⋅^oC, and the density of the solution is 1.00 g/mL.

The student uses the results to estimate the enthalpy of combustion of magnesium:

Mg (s) + O₂ (g) → MgO (s)

The enthalpy of formation of liquid water is known:

ΔH_f^∘(H₂O (l)) = −285.8 kJ/mol

Determine the heat released when 0.500 g of magnesium reacts with HCl. Include the appropriate sign.

Calculate the enthalpy change per mole of Mg for Reaction 1.

Using Hess’s Law, estimate the enthalpy change for the combustion of magnesium. Justify the signs of each term in your calculation.