Calorimetry

This question is about enthalpy changes of reactions involving hydrocarbons.

A student determines the enthalpy change of combustion, _cH, of heptane, C₇H₁₆, using the method outlined below.

• Add 150 g of water to a beaker and measure its temperature.

• Weigh a spirit burner containing heptane and use it to heat the water.

• Extinguish the flame and record the maximum temperature reached by the water.

• Reweigh the spirit burner.

The temperature of the water increased by 10.5 °C.

The spirit burner decreased in mass by 0.133 g.

Use the student’s results to determine the enthalpy change of combustion of heptane, _cH (C₇H₁₆), in kJ mol^–1.

_cH (C₇H₁₆) =........ kJ mol^–1

Nonane, C₉H₂₀, can be broken down by heat to form pentane, C₅H₁₂, and ethene, C₂H₄.

The enthalpy changes of combustion of C₉H₂₀(g) and C₂H₄(g) are shown in the table below.

Use H in Reaction 1 and the enthalpy changes of combustion in the table to determine the enthalpy change of combustion of C₅H₁₂(g).

∆_cH (C₅H₁₂(g)) =........ kJ mol^–1

Ethene can be produced from ethanol, as shown in Reaction 2 below.

Predict the sign of the entropy change, S, for Reaction 2.

Explain your reasoning.

These questions are from different areas of chemistry.

This question is about two salts of rubidium (atomic number 37): RbClO₃ and RbClO₄.

The oxidation number of chlorine is different in the two rubidium salts, RbClO₃ and RbClO₄.

What is the name of RbClO₄?

A student carries out an experiment to determine the enthalpy change of solution of RbClO₃ using the method below.

• A 2.00 g sample of solid RbClO₃ is added to water in a well-insulated container. The initial temperature is 23.0 °C.

• The mixture is stirred until all the RbClO₃ has dissolved. The final temperature is 21.5 °C. The final solution has a mass of 102 g.

Determine the enthalpy change of solution, _solH, of RbClO₃ in kJ mol^–1.

Assume that the specific heat capacity of the solution is the same as that of pure water.

_solH (RbClO₃) =........ kJ mol^–1

A student investigates the rate of a reaction that is 1st order with respect to hydrochloric acid, HC l (aq).

• The student carries out a reaction using 0.680 mol dm^–3 HCl(aq). The initial rate is

• The student dilutes a different sample of 0.680 mol dm^–3 HCl(aq) with water. The pH of this diluted acid is 1.50.

• The student repeats the reaction using the same volume of this diluted acid.

Determine the initial rate using this diluted acid.

initial rate =........ mol dm^–3 s^–1

This question is about energy changes.

Magnesium reacts with aqueous silver nitrate, AgNO₃(aq) as shown below.

A student adds an excess of magnesium to 100.0 cm³ of 0.400 mol dm^–3 AgNO₃(aq). The initial temperature is 20.0 °C.

Determine the maximum temperature reached in this reaction. Give your answer to 3 significant figures.

Assume that the specific heat capacity and density of the solution are the same as for water, and that there are no heat losses.

The student wants to repeat the experiment, but there is not enough AgNO₃(aq) left to use another 100.0 cm³ portion.

The student decides to modify the method by adding an excess of magnesium to 50.0 cm³ of 0.400 mol dm^–3 AgNO₃(aq).

Predict, with reasons, how this modification would affect the maximum temperature reached. Assume that there are no heat losses.

Nitric acid is manufactured from ammonia in a multi-stage process. The equation for the first stage in this process is shown in Reaction 17.1.

Some standard enthalpy changes of formation are shown in the table.

Compound	f H / kJ mol–1
NH3(g)	–46
H2O(l)	–286

Explain the term enthalpy change of formation.

Calculate the standard enthalpy change of formation, _fH, of NO(g).

_fH of NO(g) =........ kJ mol^–1

Carbon disulfide, CS₂, reacts with dinitrogen oxide, N₂O, as shown in Reaction 17.2.

Standard entropies, S, are shown in the table.

Substance	CS2(l)	N2O(g)	S8(s)	CO2(g)	N2(g)
S/ J K–1 mol–1	151	220	256	214	192

Explain the term entropy.

The free energy change, G, of Reaction 17.2 is –2672 kJ mol^–1 at 25 °C.

Calculate the enthalpy change, H, of Reaction 17.2, in kJ mol^–1.

H =........ kJ mol^–1

A student concludes that Reaction 17.2 is feasible at all temperatures.

Explain whether the student is correct or not.

This question is about energy changes.

* A student plans to determine the enthalpy change of hydration of calcium ions.

The student finds the information below from data tables.

The student carries out an experiment to find the enthalpy change of solution of calcium chloride.

Student’s method:

• Weigh a bottle containing calcium chloride and weigh a polystyrene cup.

• Add water from a measuring cylinder to the polystyrene cup and measure its temperature.

• Add the calcium chloride, stir the mixture, and measure the maximum temperature of the final solution.

• Weigh the empty bottle and weigh the polystyrene cup with the final solution.

Mass readings

Temperature readings

Calculate the enthalpy change of solution of calcium chloride and determine the enthalpy change of hydration of calcium ions.

Show your working, including an energy cycle linking the energy changes.

Assume that the density and specific heat capacity, c, of the solution are the same as for

water.

Internal combustion engines have historically used fuels obtained from crude oil as a source of power.

The environmental effects of fossil fuel use can be reduced by blending petrol with biofuels such as ethanol.

A fuel is being developed using a 1:1 molar ratio of octane and ethanol.

Write the equation for the complete combustion of this fuel.

Calculate the energy released, in kJ, by the complete combustion of 8.00 kg of this fuel.

_cH(C₈H₁₈) = –5470 kJ mol^–1; ∆_cH(C₂H₅OH) = –1367 kJ mol^–1.

This question is about acids, bases and buffers.

Sodium hydroxide, NaOH, is a strong base.

Calculate the pH of 0.140 mol dm^–3 NaOH(aq) at 298 K.

Give your answer to 2 decimal places.

Sulfuric acid reacts with sodium hydroxide as shown in Equation 22.1. Equation 22.1 H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l) H₁ This is a neutralisation reaction.

A student carries out an experiment to determine the enthalpy change H₁ and uses this value to deduce the enthalpy change of neutralisation, _neutH.

The student measures out two solutions:

• 25.0 cm³ of 1.60 mol dm^–3 H₂SO₄(aq)

• 55.0 cm³ of 1.50 mol dm^–3 NaOH(aq) (an excess).

The temperature of each solution is the same.

The student mixes the two solutions. The temperature increases by 13.0 °C.

Show that NaOH is in excess.

Calculate the enthalpy change, H₁, for Equation 22.1, and deduce the value for the enthalpy change of neutralisation, ∆_neutH, in kJ mol^–1.

Assume that the densities of all solutions and the specific heat capacity, c, of the reaction mixture are the same as for water.

enthalpy change, H₁ =........ kJ mol^–1

enthalpy change of neutralisation, _neutH =........ kJ mol^–1

The student repeats the experiment using 50.0 cm³ of 1.60 mol dm^–3 H₂SO₄ and 110.0 cm³ of 1.50 mol dm^–3 NaOH.

Predict the increase in temperature.

Explain your reasoning.

* Nitrous acid, HNO₂, is a weak Brønsted–Lowry acid with a pK_a value of 3.34 at room temperature.

HNO₂ can be prepared by reacting N₂O₃ with water. HNO₂ is the only product.

A chemist makes up a buffer solution by the following method.

Step 1 The chemist weighs a sample of N₂O₃. Water is then added to form 100 cm³ of 0.500 mol dm^–3 HNO₂(aq).

Step 2 The chemist adds 100 cm³ of 0.150 mol dm^–3 NaOH(aq) to the 100 cm³ solution of 0.500 mol dm^–3 solution of HNO₂(aq). The resulting solution is made up to 1.00 dm³.

Explain why a buffer solution forms in Step 2. Determine the pH of this buffer solution and the mass of N₂O₃ that was used in Step 1.

Propane, C₃H₈, (boiling point −42 °C) is used as ‘camping gas’. A student plans to determine the enthalpy change of combustion of propane, _cH (C₃H₈), by two methods.

The student first carries out an experiment using the apparatus below.

Determine the enthalpy change of combustion of propane, in kJ mol⁻¹.

Give your answer to 3 significant figures.

_cH (C₃H₈) =........ kJ mol⁻¹

The student finds that the experimental enthalpy change _cH (C₃H₈) is much less exothermic than the accurate standard enthalpy change _cH (C₃H₈) in databases.

One reason could be that the student’s experiment had not been carried out under standard conditions.

Suggest two other reasons for this difference in enthalpy change.

This question is about copper(II) sulfate, CuSO₄, and sodium thiosulfate, Na₂S₂O₃.

* Calculate ∆_solH(CuSO₄(s)) for reaction 5.2 and determine the enthalpy change of reaction 5.1, ∆_rH.

Assume that the specific heat capacity, c, of the solution is the same as for water.

Show your working, including an energy cycle linking the enthalpy changes.

The standard enthalpy change of reaction, ∆_r , and the standard free energy change, ∆ for converting anhydrous sodium thiosulfate to hydrated sodium thiosulfate are shown below.

Standard entropies are given in the table.

Determine the standard entropy, , of anhydrous sodium thiosulfate, Na₂S₂O₃(s).

Give your answer to 3 significant figures.

Sodium thiosulfate contains the thiosulfate ion, S₂O₃²⁻. The displayed formula of S₂O₃²⁻ can be shown as below.

Predict the O–S–S bond angle and name of the shape of the thiosulfate ion.

Bond angle........

Name of shape........

In some of its reactions, the thiosulfate ion forms the tetrathionate ion, S₄O₆²⁻.

The S₄O₆²⁻ ion is a ‘dimer’ of S₂O₃²⁻.

Draw a displayed formula for the S₄O₆²⁻ ion.

A student carries out two experiments in the laboratory based on succinic acid (butanedioic acid), (CH₂COOH)₂.

Succinic acid is esterified by ethanol, C₂H₅OH, in the presence of an acid catalyst to form an equilibrium mixture.

The equilibrium constant, K_c, for this equilibrium can be calculated using the amounts, in moles, of the components in the equilibrium mixture, using expression 5.1.

A student carries out an experiment to determine the value of K_c for this equilibrium.

• The student mixes together 0.0500 mol of succinic acid and 0.150 mol of ethanol, with a small amount of an acid catalyst.

• The mixture is allowed to reach equilibrium.

• The student determines that 0.0200 mol of succinic acid are present in the equilibrium mixture.

Which technique could be used to determine the equilibrium amount of succinic acid?

Write the equation for the equilibrium reaction that takes place.

Draw the skeletal formula of the ester present in the equilibrium mixture.

K_c is the equilibrium constant in terms of equilibrium concentrations.

Why can expression 5.1 be used to calculate K_c for this equilibrium?

Calculate the value of K_c for this reaction.

Show your working.