Energetics: AQA A-Level Chemistry past-paper questions

07.1 Define enthalpy change.

07.2 The enthalpy of combustion of ethanol is –1371 kJ mol⁻¹

A sample of ethanol is burned completely in air to raise the temperature of

55.0 g of water by 46.2 °C

Calculate the mass of ethanol in the sample.

Assume there is no heat loss.

Specific heat capacity of water, c = 4.18 J K⁻¹ g⁻¹

Mass of ethanol g

An alkane burns completely in oxygen. The equation for this reaction is shown.

3n + 1

C_nH_2n+2 + ₂ → n CO₂ + (n+1)H₂O

2 O

Table 2 gives some mean bond enthalpies.

Table 2

07.3 Use Table 2 to show that an expression for the sum of the bond enthalpies for all the bonds in the products is 2536n + 926

07.4 An expression for the sum of the bond enthalpies for all the bonds in the reactants is

1916n + 724

A calculated value for the enthalpy of combustion of an alkane is –3302 kJ mol⁻¹

Use this enthalpy change and the expressions to calculate a value of n for this alkane.

Use your value of n to deduce the molecular formula of the alkane.

Value of n

Molecular formula

01.1 Give the equation, including state symbols, for the process that occurs when the enthalpy of solution of potassium hydroxide is determined.

01.2 A student does an experiment to determine a value for the enthalpy of solution of potassium hydroxide.

The student uses this method.

• Add some distilled water to a measuring cylinder and record the volume.
• Pour the water into a polystyrene cup.
• Record the temperature of the water in the cup.
• Add 3.00 g of solid KOH to the water in the cup.
• Stir the solution and record the maximum temperature reached.

Table 1 shows the student’s results.

Table 1

The student calculated the enthalpy of solution of potassium hydroxide to be –51.3 kJ mol⁻¹

The specific heat capacity of the solution, c = 4.18 J K⁻¹ g⁻¹

Assume the density of the solution = 1.00 g cm⁻³

Calculate the volume of water that the student added to the cup.

Volume of water Units

01.3 Table 2 shows some enthalpy change data.

Table 2

Use the data in Table 2 to calculate the enthalpy of solution for potassium hydroxide.

Enthalpy of solution kJ mol⁻¹

01.4 Suggest why the value calculated in Question 01.3 is different from the experimental value given in Question 01.2.

(a) Define standard enthalpy of formation.

(b) Silver nitrate(V) is formed when silver nitrate(III) undergoes thermal decomposition.

2 AgNO₂(s) → Ag(s) + AgNO₃(s) + NO(g)  ΔH = +56.2 kJ mol⁻¹

The standard enthalpy of formation of AgNO₃(s) is −123.0 kJ mol⁻¹

The standard enthalpy of formation of NO(g) is +90.4 kJ mol⁻¹

Determine the standard enthalpy of formation of AgNO₂(s)

Standard enthalpy of formation kJ mol⁻¹

(c) Suggest why the enthalpy change for the thermal decomposition of solid silver nitrate(III) is difficult to determine experimentally.

Silver nitrate(V) solution reacts with solid sodium chloride.

AgNO₃(aq) + NaCl(s) → AgCl(s) + NaNO₃(aq)

A student does an experiment to determine the enthalpy change for this reaction.

The student follows this method:

1.  Measure out 50 cm³ of 0.10 mol dm⁻³ aqueous silver nitrate(V) using a clean, dry measuring cylinder.

2.  Pour the silver nitrate(V) solution into a glass beaker.

3.  Weigh out 2.00 g of solid sodium chloride (an excess) using a weighing boat and tip the solid into the silver nitrate(V) solution. Reweigh the weighing boat to determine the mass of sodium chloride added.

4.  Add a lid to the beaker that has two small holes for a stirring rod and for a thermometer.

5.  Stir the mixture with a plastic stirring rod whilst recording the temperature with a thermometer.

6.  Record the maximum temperature reached.

(d) Identify three aspects of this method which could cause inaccurate results.

Describe how the student could improve these three aspects of the method to obtain more accurate results.

Inaccuracy 1

Improvement 1

Inaccuracy 2

Improvement 2

Inaccuracy 3

Improvement 3

(a) Limonene is found in the skin of citrus fruits.

The figure below shows a diagram of the apparatus used in an experiment to determine a value for the enthalpy of combustion of limonene.

When 1.31 g of limonene are burned, the temperature of the 60.0 g of water in the copper calorimeter increases by 52.1 °C

The specific heat capacity of water is 4.18 J K⁻¹ g⁻¹

Calculate a value for the enthalpy of combustion, in kJ mol⁻¹, of limonene (C₁₀H₁₆).

Enthalpy of combustion kJ mol⁻¹

(b) The table below shows values, obtained by different methods, for the enthalpy of combustion of a different liquid hydrocarbon.

Suggest reasons for the differences between the values obtained by each of Methods 2 and 3, and the value obtained by Method 1 in the table above.

(a) Complete the equation for the hydrolysis of the triester in palm oil under acidic conditions.

(b) Palmitic acid burns in air.

In a calorimetry experiment, combustion of 387 mg of palmitic acid increases the temperature of 0.150 kg of water from 23.9 °C to 37.5 °C

Calculate a value, in kJ mol⁻¹, for the enthalpy of combustion of palmitic acid in this experiment.

Give your answer to the appropriate number of significant figures.

The specific heat capacity of water is 4.18 J K⁻¹ g⁻¹

Enthalpy of combustion kJ mol⁻¹

(c) State how the value calculated in part (b) is likely to differ from data book values.

Give one reason, other than heat loss, for this difference.

Difference

Reason

(d) A sample of a different biofuel, made from sewage sludge, is found to contain 37.08% carbon, 5.15% hydrogen and 24.72% oxygen by mass.

The rest of the sample is sulfur.

Calculate the empirical formula of this biofuel.

Empirical formula

(e) Complete combustion of the biofuel made from sewage sludge produces the greenhouse gas carbon dioxide.

Suggest one other possible environmental problem with the complete combustion of this biofuel.

State the formula of the pollutant responsible for this problem.

Environmental problem

Formula

(f) Ethanol is a biofuel that can be produced by the fermentation of glucose.

C₆H₁₂O₆ → 2 C₂H₅OH + 2 CO₂

Glucose has the structural formula shown.

Table 1 shows some mean bond enthalpy values.

Table 1

Use the equation and the data in Table 1 to calculate an approximate value of ΔH for the fermentation of glucose. For this calculation you should assume that all the substances are in the gaseous state.

ΔH kJ mol⁻¹

(g) The carbon dioxide produced from fermentation can be reacted with steam to make more ethanol.

The equation for this reaction is

2 CO₂(g) + 3 H₂O(g) → C₂H₅OH(g) + 3 O₂(g)

Table 2 shows some standard enthalpies of formation.

Table 2

Use the data in Table 2 to calculate a standard enthalpy change value for this reaction.

Standard enthalpy change kJ mol⁻¹

(a) The student finds that when all the alcohol is burned, the temperature of the water increases from 18.9 °C to 78.1 °C

Calculate the enthalpy of combustion, in kJ mol⁻¹, for the alcohol.

The specific heat capacity of water, c = 4.18 J g⁻¹ K⁻¹

Enthalpy of combustion kJ mol⁻¹

(b) Table 1 shows the enthalpies of combustion of some alcohols.

Table 1

Explain how your answer to part (a) suggests that the alcohol is butan-1-ol.

(If you have been unable to obtain an answer for part (a), assume that the answer is –2120 kJ mol⁻¹)

(c) The equation for the complete combustion of gaseous pentan-1-ol is shown.

CH₃(CH₂)₃CH₂OH(g) + 7½O₂(g) → 5 CO₂(g) + 6 H₂O(g)  ΔH = –3388 kJ mol⁻¹

Table 2 shows some bond enthalpy data.

Table 2

Use data from Table 2 to calculate a value for the mean C–C bond enthalpy in pentan-1-ol.

C–C bond enthalpy kJ mol⁻¹

(d) The energy stored in fuels can be compared using energy density values measured in kJ dm⁻³

Calculate the energy density of butan-1-ol.

enthalpy of combustion of butan-1-ol = –2676 kJ mol⁻¹

density of butan-1-ol = 0.810 kg dm⁻³

relative molecular mass (M_r) of butan-1-ol = 74.0

Energy density kJ dm⁻³

(a) The ethanol produced by fermentation of glucose may be regarded as a carbon-neutral fuel.

Justify this statement. Include the relevant chemical equations in your answer.

Coffee beans from South America are exported to Europe in an outer layer called silverskin.

The waste silverskin can be fermented to produce a solution containing propanone, ethanol and butan-1-ol.

(b) Suggest why ethanol produced in Europe using silverskin from South America is less likely to be carbon-neutral than ethanol produced from crops grown in Europe.

(c) Table 1 shows the enthalpies of combustion of the three fuels from the fermentation of silverskin.

One way to measure a fuel’s environmental impact is to measure the amount of energy released per mole of CO₂ produced.

Complete Table 1.

Use your answers to deduce the fuel with the lowest environmental impact by this measure.

(d) A student investigated the combustion of propanone (C₃H₆O) using calorimetry.

A copper calorimeter containing water was heated by the complete combustion of some propanone. The student did not record the final temperature of the water.

Table 2 shows the student’s results.

Use the results in Table 2 to calculate a value for final temperature of the water in the experiment.

Assume that no heat was lost in the experiment and that the heat capacity of the calorimeter is negligible.

For propanone, enthalpy of combustion = –1786 kJ mol⁻¹

For water, specific heat capacity = 4.18 J g⁻¹ K⁻¹

Final temperature of water °C

(e) Butan-1-ol can be added to petrol for cars.

An equation for the complete combustion of gaseous butan-1-ol is shown.

C₄H₉OH(g) + 6 O₂(g) → 4 CO₂(g) + 5 H₂O(g) ΔH = –2504 kJ mol⁻¹

Table 3 shows some mean bond enthalpy data.

Use these data to calculate a value for the mean C–C bond enthalpy in gaseous butan-1-ol.

C–C bond enthalpy kJ mol⁻¹

Butan-1-ol can be manufactured by reacting steam with but-1-ene in the presence of the catalyst, concentrated sulfuric acid.

In the first part of this process, but-1-ene reacts with concentrated sulfuric acid to form compounds W and X.

Butan-1-ol is then made from compound W.

(f) Name and outline a mechanism to show the conversion of but-1-ene into compound W in the first part of this process.

Name of mechanism

Outline of mechanism

(g) There is a very low yield of butan-1-ol from but-1-ene in this manufacturing process.

Explain why.

(a) Design a table to record all the readings necessary to determine an experimental value for the enthalpy of combustion for heptane in this experiment.

(b) The student considered using a glass beaker on a tripod and gauze instead of the clamped copper calorimeter.

Suggest two disadvantages of using a glass beaker on a tripod and gauze.

Disadvantage 1

Disadvantage 2

(c) Suggest two reasons why the value of enthalpy of combustion from this experiment is less exothermic than a data book value.

Reason 1

Reason 2

(d) Suggest one addition to this apparatus that would improve the accuracy of the enthalpy value obtained.

(a) Define the term enthalpy change.

(b) Propane undergoes complete combustion.

C₃H₈(g) + 5 O₂(g) → 3 CO₂(g) + 4 H₂O(l) ΔH = –2046 kJ mol⁻¹

The table below shows some bond enthalpy data.

The bond enthalpy for O=O is 496 kJ mol⁻¹

For H₂O(l) → H₂O(g) ΔH = +41 kJ mol⁻¹

Use these data to calculate a value for the C–C bond enthalpy in propane.

C–C bond enthalpy kJ mol⁻¹

(c) Explain why the value given for the O=O bond enthalpy in part (b) is not a mean value.

(a) State the meaning of the term standard enthalpy of combustion.

(b) A student does an experiment to determine the enthalpy of combustion of propan-1-ol (CH₃CH₂CH₂OH, M_r = 60.0).

Combustion of 0.497 g of propan-1-ol increases the temperature of 150 g of water from 21.2 °C to 35.1 °C

Calculate a value, in kJ mol⁻¹, for the enthalpy of combustion of propan-1-ol in this experiment.

The specific heat capacity of water is 4.18 J K⁻¹ g⁻¹

Enthalpy of combustion kJ mol⁻¹

(c) The enthalpy of combustion determined experimentally is less exothermic than that calculated using enthalpies of formation.

Give one possible reason for this, other than heat loss.

(a) Define the term standard enthalpy of formation.

(b) The equation for a reaction used to manufacture 1-chloropropane is

3 CH₃CH₂CH₂OH(l) + PCl₃(l) → 3 CH₃CH₂CH₂Cl(l) + H₃PO₃(s)

The enthalpy change for this reaction, ΔH, is –114 kJ mol⁻¹

The table contains some standard enthalpy of formation data.

Calculate a value for the standard enthalpy of formation of propan-1-ol using the enthalpy change for the reaction and data from the table.

Standard enthalpy of formation kJ mol⁻¹

(c) 1-chloropropane can also be produced by the reaction between propane and chlorine in the presence of ultraviolet light.

State why ultraviolet light is needed for this reaction to occur.

Give an equation for each propagation step in the formation of 1-chloropropane from propane.

Why ultraviolet light is needed

Propagation step 1

Propagation step 2

(d) The C–Cl bond in 1-chloropropane is polar because carbon and chlorine have different electronegativities.

Define the term electronegativity.

(e) Ammonia reacts with 1-chloropropane to form propylamine.

Name and outline the mechanism for this reaction.

Name of mechanism

Outline of mechanism

(a) The heat energy released in the calorimeter, q = CcalΔT

Calculate the heat capacity (Ccal) in kJ K⁻¹

Ccal kJ K⁻¹

(b) When the experiment is repeated with 2.00 g of octane (M_r = 114.0) the temperature change recorded is 12.2 °C

Calculate the heat change, in kJ mol⁻¹, for octane in this combustion reaction.

If you were unable to calculate a value for Ccal in part (a), use 6.52 kJ K⁻¹ (this is not the correct value).

Heat change kJ mol⁻¹

(c) State why the heat change calculated from the bomb calorimeter experiment is not an enthalpy change.

(d) The thermometer used to measure the temperature change of 12.2 °C in part (b) has an uncertainty of ± 0.1 °C in each reading.

Calculate the percentage uncertainty in this use of the thermometer.

Suggest one change to this experiment that decreases the percentage uncertainty while using the same thermometer.

Percentage uncertainty

Change

(a) State the meaning of the term enthalpy change as applied to a chemical reaction.

(b) A student determines the enthalpy change for the reaction between calcium carbonate and hydrochloric acid.

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

The student follows this method:

• measure out 50 cm³ of 1.00 mol dm⁻³ aqueous hydrochloric acid using a measuring cylinder and pour the acid into a 100 cm³ glass beaker
• weigh out 2.50 g of solid calcium carbonate on a watch glass and tip the solid into the acid
• stir the mixture with a thermometer
• record the maximum temperature reached.

The student uses the data to determine a value for the enthalpy change.

Explain how the experimental method and use of apparatus can be improved to provide more accurate data.

Describe how this data from the improved method can be used to determine an accurate value for the temperature change.

(c) In a different experiment 50.0 cm³ of 0.500 mol dm⁻³ aqueous hydrochloric acid are reacted with 50.0 cm³ of 0.500 mol dm⁻³ aqueous sodium hydroxide.

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l) ΔH = –57.1 kJ mol⁻¹

The initial temperature of each solution is 18.5 °C

Calculate the maximum final temperature of the reaction mixture.

Assume that the specific heat capacity of the reaction mixture, c = 4.18 J K⁻¹ g⁻¹

Assume that the density of the reaction mixture = 1.00 g cm⁻³

Final temperature °C

(d) Suggest how, without changing the apparatus, the experiment in part (c) could be improved to reduce the percentage uncertainty in the temperature change.

02.1 Write an equation, including state symbols, for the reaction with an enthalpy change equal to the enthalpy of formation for iron(III) oxide.

02.2 Table 2 contains some standard enthalpy of formation data.

Table 2

Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g) ΔH = −19 kJ mol⁻¹

Use these data and the equation for the reaction of iron(III) oxide with carbon monoxide to calculate a value for the standard enthalpy of formation for carbon dioxide.

Show your working.

ΔfH^⦵ kJ mol⁻¹

02.3 Some enthalpy data are given in Table 3.

Table 3

Use the data from Table 3 to calculate the bond enthalpy for N–H in ammonia.

N–H bond enthalpy kJ mol⁻¹

02.4 Give one reason why the bond enthalpy that you calculated in Question 2.3 is different from the mean bond enthalpy quoted in a data book (388 kJ mol⁻¹).

03.1 Show by calculation which reactant was in excess.

Use the data to calculate the experimental value for enthalpy of reaction in kJ mol⁻¹(Assume that the specific heat capacity of the solution is

4.18 J K⁻¹g⁻¹and the density of the solution is 1.00 g cm⁻³).

Reactant in excess

Enthalpy of reaction kJ mol⁻¹

03.2 Another student used the same method and obtained a value for the enthalpy of reaction of −142 kJ mol⁻¹

A data book value for the enthalpy of reaction is −310 kJ mol⁻¹

Suggest the most likely reason for the large difference between the student’s experimental value and the data book value.

03.3 Suggest how the students’ method, and the analysis of the results, could be improved in order to determine a more accurate value for the enthalpy of reaction.

Justify your suggestions.

Do not refer to the precision of the measuring equipment. Do not change the amounts or the concentration of the chemicals.