Energetics (AQA A Level Chemistry)

Exam Questions

3 hours45 questions
1a2 marks

During chemical reactions, enthalpy changes occur as bonds are broken and formed.

i)
Thermal energy is needed to overcome the attractive forces between atoms. In terms of thermal energy, name the process where bonds are broken.

ii)
When bonds are formed, thermal energy is released to the surroundings. In terms of thermal energy, name the process where bonds are made.
1b1 mark

Is the reaction exothermic or endothermic if more thermal energy is released making new bonds than taken in breaking bonds?

1c5 marks

The energy level diagram for an endothermic reaction is shown in Figure 1.

Figure 1

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Label the energy level diagram shown in Figure 1.

1d5 marks

An element X undergoes complete combustion according to the following equation.

X + O2 → XO2

The enthalpy change, ΔH, and activation energy, Ea, for this reaction are -520 kJ mol-1 and +630 kJ mol-1 respectively.

Draw the energy level diagram for this reaction.

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2a3 marks

Define the term enthalpy of neutralisation.

2b6 marks
i)
Define the term enthalpy of formation.

ii)
State the standard conditions.

2c3 marks

Identify the following reactions as:

  • Enthalpy of reaction / ΔHr
  • Enthalpy of formation / ΔHf
  • Enthalpy of combustion /  ΔHc
  • Enthalpy of neutralisation / ΔHneut 

i)
C2H5OH (l) + O2 (g) → CO2 (g) + H2O (l)

ii)
CaSO3 (s) CaO (s) + SO2 (g)

iii)
S (s) + O2 (g) SO2 (g)
2d1 mark

Suggest why the enthalpy of formation, ΔHf, of liquid iron is +18.3 kJ mol-1 and not 0 kJ mol-1.

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3a3 marks

The equipment set up in Figure 1 is used to measure the enthalpy change for a reaction.

Figure 1

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Suggest suitable materials for the lid and cup. Justify your choices.

3b2 marks

The equation that is commonly used in calorimetry questions is Q = mcΔT. Complete Table 1 with information about each term in the equation.

Table 1

Term 

Represents 

Units 

Q

Thermal energy transferred 

 

m

   

c

 

J g-1 K-1 or J g-1 bold degreeC-1 

ΔT

   
3c3 marks

In a calorimetry experiment, 2.50 g of ethanol is burnt in excess oxygen. The energy released during the combustion is absorbed by 250 cm3 of water, which increases in temperature from 20 degreeC to 71 degreeC

The specific heat capacity of water is 4.18 J g-1 degreeC-1  

i)
Calculate the energy released in this experiment using the equation Q = mcΔT.

ii)
Convert your answer to part (i) into kJ. Give your answer to 3 significant figures.
3d1 mark

The equipment set up from part (a), Figure 1 is repeated here for your use.

Figure 1

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The actual amount of energy that should be released by 2.50 g of ethanol is 74.3 kJ.

Why is your answer to part (c) lower than the actual value? 

If you did not calculate an answer to part (c) then you should use a value of 51.3 kJ, thisis not the correct answer.

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4a1 mark

State Hess’s Law.

4b2 marks

Bond energies cannot be found directly. Hess cycles use the enthalpies of atomisation with either formation or combustion to find the average bond energy.

Define enthalpy of atomisation, ΔHat.

4c7 marks

Copper carbonate undergoes thermal decomposition according to the following equation.

CuCO3 (s) → CuO (s) + CO2 (g)

i)
Complete the Hess cycle diagram in Figure 1 by adding the correct chemicals to each box and labelling the arrows. One box has been filled and one arrow has been labelled for you.

Figure 1

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ii)
Use the information in Table 1 to calculate the enthalpy of reaction for the thermal decomposition of copper carbonate.

Table 1

Molecule

Enthalpy of formation / kJ mol-1 

CuCO3 (s)

-595

CuO (s)

−155

CO2 (g)

−394

4d3 marks

The Hess cycle to calculate the average bond enthalpy of the S=O bond is shown in Figure 2.

Figure 2

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i)
Use the information in Table 2 to calculate ΔH.

Table 2

Molecule

Energy / kJ mol-1 

ΔHf SO2 (g)

−297

ΔHat S (g)

+279

ΔHat ½ O2 (g)

+249

ii)
Sulfur dioxide contains two S=O bonds. Use your answer from part (i) to calculate the bond enthalpy of one S=O bond.

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5a1 mark

Define bond dissociation energy.

5b1 mark

Write the equation to calculate the enthalpy of reaction using bond energies.

5c2 marks

Using displayed formulae, write the equation for the reaction of ethene with water to form ethanol.

5d4 marks

Use the information in Table 1 to calculate the enthalpy of reaction for ethene reacting with water to form ethanol. 

You should use your mathematical equation from part (b) and your displayed reaction equation from part (c).

Table 1

Bond 

Bond enthalpy / kJ mol-1 

C-C

+347

C=C

+612

C-H

+413

O-H

+464

C-O

+358

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1a1 mark

Alkanes can be used as fuels in internal combustion engines. When sufficient oxygen is present, they undergo complete combustion reactions.

Write an equation for the complete combustion of butane.

1b3 marks

Define the term standard enthalpy of combustion, ΔcHϴ

1c5 marks

Table 1 below contains bond enthalpy data for the reaction shown in part (a). 

Table 1

 

C-C

C-H

O=O

C=O

O-H

Mean bond enthalpy (kJ mol-1)

348

412

496

805

463

i)
Using the data in Table 1 and the equation in part (a), calculate the enthalpy change of combustion of butane.

ii)
State whether this reaction is endothermic or exothermic. Explain your answer.
1d2 marks

The values given in Table 1 in part (c) are mean bond enthalpies.

State the meaning of the term mean bond enthalpies. 

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2a3 marks

Define the term standard enthalpy of formation, ΔfHϴ.

2b1 mark

Ammonia is a colourless gas which has a very pungent smell. It also has a number of important chemical applications. 

Ammonia reacts with hydrogen chloride, to produce solid ammonium chloride. 

Write the equation, including state symbols, for this reaction.

2c3 marks

Some enthalpy of formation data are shown below in Table 1

Table 1

 

NH3

HCl

NH4Cl

Enthalpy of formation, ΔfHϴ (kJ mol-1)

-46

-92

-314

Using the data provided in Table 1, calculate the enthalpy change for the reaction between ammonia and hydrogen chloride in part (b). 
Show your working. 

2d1 mark

Ammonia is produced via the Haber process, where hydrogen and nitrogen gas are reacted together. 

State why the enthalpy of formation of hydrogen would be zero. 

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3a3 marks

A teacher instructs a class to complete a calorimetry practical, to calculate the enthalpy change that occurs when hydrochloric acid and sodium hydroxide react together. 

Each student was given roughly 60 cm3 of 0.35 mol dm-3 hydrochloric acid, roughly 60 cm3 of 0.35 mol dm-3 sodium hydroxide, a polystyrene cup and access to all standard laboratory equipment. 

Draw a diagram to demonstrate the practical set up that the students would need to use to determine the enthalpy change during this neutralisation reaction, and state the key measurements that the students would have to make. 

3b5 marks

The students then completed the practical from part (a), using their own method and measurements that they had chosen.

One student found that when they reacted 35.0 cm3 of the hydrochloric acid with 35.0 cm3 of the sodium hydroxide, the temperature rose from 19.6 degreeC to 22.3 degreeC. 

Calculate the enthalpy change for this reaction in kJ mol-1

Assume that both solutions have a density of 1.00 g cm-3 and a specific heat capacity of 4.18 J K-1 g-1

3c1 mark

Explain why the value that you have calculated for the students’ practical in part (b), might be different from the correct value given in a data book. 

3d1 mark

State how the students' practical could be improved to allow the students to calculate a more accurate value which is closer to the correct value given in data books. 

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4a1 mark

Propane gas is commonly used as a fuel for outdoor cooking. It can be produced in a number of ways, including from the addition reaction of propyne gas with hydrogen. Propyne has the formula CHCCH3 and includes a triple bond. 

Write an equation for the formation of propane from propyne. 

4b3 marks

The table below contains some key enthalpy of combustion data. 

Table 1

 

Propyne

Hydrogen

Propane

Enthalpy of combustion, ΔcHϴ (kJ mol-1)

-1940

-285.8

-2220

Calculate the enthalpy change in kJ mol-1 for the reaction in part (a) using the data provided in Table 1.

Show all working. 

4c3 marks

Table 2 below has some enthalpy data for a different chemical reaction. Hydrazine, N2H4 can react with hydrogen peroxide in an exothermic reaction, as shown below. 

N2H4 (g) + 2H2O2 → N2 (g) + 4H2O (g)     ΔrHϴ = -789 kJ mol-1 

The structure of hydrazine is shown in Figure 1

Figure 1

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Table 2

Chemical bond

ΔHϴ (kJ mol-1)

N-N

+163

N≡N

+944

O-H

+463

O-O

+146

Using the reaction equation and the data in the table above, calculate the value of the N-H bond in hydrazine. 

4d1 mark

Often when a bond enthalpy is calculated, it is a different value to the one which is quoted in a data book. State a reason for this.

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5a5 marks

Alcohols are often used as fuels. They can be used to fill spirit burners, which can be used in calorimetry practicals by students.

A teacher set a class a calorimetry practical to calculate the enthalpy of combustion of ethanol, propan-1-ol and butan-1-ol. The set up that the student used for the experiment is shown in Figure 1

Figure 1

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Table 1 shows the change in mass when using the ethanol spirit burner. 

Table 1

Mass of spirit burner before (g)

Mass of spirit burner after (g)

197.6

195.1

The temperature increased by 34.5 degreeC during the reaction.
Using the students results calculate the enthalpy of combustion for ethanol in this experiment.
The specific heat capacity of the water is 4.18 J K-1.

5b2 marks

Suggest two reasons, excluding heat transfer, why the value which you have calculated from the students’ data in part (a), might not be the same as a value provided in a data book.

5c1 mark

The student repeated the experiment in part (a) with propan-1-ol and butan-1-ol. 

Write an equation for the complete combustion of one mole of butan-1-ol. 

5d3 marks

Table 2, shown below, contains enthalpy data for butanol, carbon dioxide and water. 

Table 2

 

C4H9OH

CO2

H2O

Enthalpy of formation, ΔfHϴ (kJ mol-1)

-327.4

-393.5

-285.8

Use the data in Table 2 and your equation in part (c), to calculate the enthalpy of combustion of butan-1-ol. 

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1a1 mark

Iron and titanium are both metals with a variety of industrial uses. They are commonly found in and extracted from ores using carbon monoxide.

State why iron and titanium are not extracted directly from their ores using carbon.

1b6 marks

Strontium salts have a number of applications such as fireworks, flares, glow in the dark paint and toothpaste for sensitive teeth. The strontium required for these salts can be extracted from the ore strontia, SrO, by displacement with powdered aluminium in a vacuum.

i)
Write a balanced symbol equation, including state symbols, for the reaction of strontia with aluminium.

ii)
State the role of the aluminium in this reaction.

iii)
The standard enthalpy change for this extraction of strontium is 99.3 kJ mol-1 and the standard enthalpy of formation of aluminium oxide is -1676.7 kJ mol-1.
Use this information to calculate the standard enthalpy of formation of strontia.
1c4 marks

Vanadium is an important metal in the alloy industry. When alloyed with iron or steel, the vanadium alloy is strong and resistant to corrosion. These alloys are often used to make parts for machinery such as axles, gears and crankshafts. When combined with gallium, it forms superconductive magnets and with aluminium and titanium, it forms a very strong alloy that is used in dental implants and jet engines.

Vanadium is commonly found in different ores such as magnetite, vanadinite and patronite. The vanadium is commonly extracted from these ores by reduction and displacement. 

Vanadium can be extracted by the reduction of vanadium pentoxide, V2O5, with calcium at high temperatures, according to the following equation.

V2O5 (s) + 5 Ca (s) → 2 V (s) + 5 CaO (s) 

i)
Explain the effect that increasing the pressure will have on the reduction of vanadium.

ii)
The enthalpy of formation of vanadium pentoxide is -1560 kJ mol-1 and the standard enthalpy change for the reaction is -1615 kJ mol-1.
Use this information to calculate the enthalpy of formation of calcium oxide.
1d6 marks

Manganese is too brittle for use as a pure metal, so it is often alloyed with other metals. Manganese is used in steel to increase the strength and resistance to wear. Manganese steel (13% Mn) is extremely strong and used for railway tracks, safes and prison bars. Alloys of 1.5% manganese with aluminium are used to make drinks cans due to the improved corrosion resistance of the alloy.

Manganese is extracted from different ores by reduction with carbon monoxide.

Mn2O3 (s) + 3 CO (g) → 2 Mn (s) + 3 CO2 (g)

Although this is not a reversible reaction, an equilibrium is established in terms of the yield and rate of reaction.

Table 1 shows some enthalpy of formation data.

Table 1

 

Mn2O3 (s) 

CO (g)

Mn (s)

CO2 (g)

ΔfHϴ / kJ mol−1

−971

−111

0

−394

Use the information to explain how temperature can be altered to increase the yield of the reaction and explain the effect that this would have on the rate of reaction.

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2a2 marks

A student investigated the temperature change for the neutralisation of malonic acid, HOOCCH2COOH, and sodium hydroxide solution according to the following method.

  1. Measure 25.0 cm3 of 0.400 mol dm–3 of malonic acid into a beaker.
  2. Record the temperature every minute for three minutes.
  3. In the fourth minute, add 50.0 cm3 of 0.500 mol dm–3 sodium hydroxide solution and stir.
  4. Record the temperature every minute for eight minutes.
The student’s results are shown in Figure 1.
Figure 1 
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Explain why the student’s temperature change of 3.5 degreeC is not an accurate temperature change.

2b1 mark

Calculate the percentage uncertainty in the student’s 3.5 degreeC temperature rise using a thermometer with an uncertainty of ±0.1 degreeC.

2c5 marks

Another student completed the same investigation and recorded a maximum temperature of 22.4 degreeC. The student calculated the heat energy, q, for the reaction to be 9.719 x 10-1 kJ.

i)
Use the information in part (a) to estimate the initial temperature for this student’s investigation. 
Assume that both solutions have the same initial temperature, a density of 1.0 g cm–3 and a specific heat capacity of 4.18 J K–1 g–1

ii)
Use the information in part (a) to explain why your calculated value for the initial temperature is an estimate.
2d4 marks
i)
Write a balanced symbol equation for the neutralisation of malonic acid with sodium hydroxide solution.

ii)
Use the information in parts (a) and (c) to calculate the enthalpy change (ΔH) per mole of water formed in this reaction. Give your answer to the appropriate number of significant figures. 

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3a5 marks

The alkanes are a homologous series of hydrocarbons with the general formula CnH2n+2

Table 1 contains the enthalpy change values for the combustion of 1 g of some alkanes.

Table 1

Alkane

Methane

Ethane

Propane

Butane

Pentane

Hexane

Heptane

Octane

ΔcHϴ / kJ per g

-55.6

-52.0

-50.4

 

-48.7

-48.4

-48.2

-48.0

Plot this data on the grid in Figure 1. Use a line of best fit to estimate the enthalpy change of combustion for 1 gram of butane.

Figure 1

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3b4 marks

Write a balanced symbol equation for the complete combustion of butane.

Table 2 shows some bond enthalpy data.

Table 2

 

C-C

C-H

O-H

C=O

O=O

Bond enthalpy / kJ mol−1

351

412

463

805

496


Using your reaction equation from part (i) and the data from Table 2 to calculate a value for the enthalpy of combustion of butane.

3c2 marks

The standard enthalpy of combustion for butane is -2878 kJ mol-1.

Calculate the percentage error of the enthalpy of combustion calculated using bond enthalpy data in part (b).

(If you did not calculate a value using the bond enthalpy data in part (b), assume that the enthalpy of combustion was -2724 kJ mol-1. This is not the correct answer.)

3d3 marks

The standard enthalpy of formation for butane cannot be measured directly. 

Table 3 shows some standard enthalpy of combustion data.

Table 3

 

C4H10 (g)

C (s), graphite

H2 (g)

ΔcHϴ / kJ mol−1

-2878

-394

-286

Use the information in Table 3 to calculate the standard enthalpy of formation for butane.

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4a2 marks

In industry, ammonia solution and carbon dioxide are passed through a saturated solution of sodium chloride to produce a precipitate of sodium hydrogen carbonate in ammonium chloride solution.

Write a balanced symbol equation, including state symbols, for this reaction.

4b6 marks

A student investigates the enthalpy change that occurs when 25 cm3 of 2.00 mol dm-3 hydrochloric acid and 3.50 g sodium hydrogen carbonate react together. 

Outline the method that the student would use to determine an accurate temperature change for the reaction.

4c4 marks

The student performed the investigation described in part (b) under standard conditions and found that the temperature decreased to 12.6 degreeC.

Calculate the enthalpy change for this reaction in kJ mol-1

Assume that the solution has a density of 1.00 g cm-3 and a specific heat capacity of 4.18 J K-1 g-1.

4d6 marks

Under standard conditions, a 6.00 g sample of sodium bicarbonate, NaHCO3, undergoes thermal decomposition to form sodium carbonate, 371.5 cm3 of a gaseous product and a liquid.

The gas constant, R, = 8.31 J K-1 mol-1.

Calculate the percentage purity of the 6.00 g sample of sodium bicarbonate.

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5a2 marks

One method to determine enthalpy changes of combustion involves the use of calorimetry.

Figure 1 shows a simple calorimeter and Figure 2 shows a chamber calorimeter.

Figure 1

Figure 2

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Suggest two reasons, other than heat loss, why the chamber calorimeter will give more accurate enthalpy changes for flammable liquids.

5b2 marks

Enthalpy changes can be calculated indirectly using Hess’s Law. Figure 1 shows a scheme of reactions involved in the formation of barium chloride solution.

Figure 1

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Table 1 shows the relevant enthalpy change data.

Table 1

 

ΔH / kJ mol-1

Ba2+ (g) + 2 Cl- (g) → BaCl2 (s) 

-2018

BaCl2 (s) → BaCl2 (aq)

- 70

Ba2+ (g) → Ba2+ (aq)

- 1360

Cl- (g) → Cl- (aq)

To be calculated

Use the information on Table 1 to calculate a value for the enthalpy of hydration of the chloride ion.

5c4 marks

The enthalpy change of solution for calcium chloride (Mr 98) can be measured using calorimetry. Data suggests the value for this enthalpy change is -120 kJ mol-1.

Calculate the expected final temperature when 1.02 g of calcium chloride is dissolved in 25.0 cm3 of water at 19.5 degreeC. 

Assume that the water has a density of 1.00 g cm–3 and a specific heat capacity of 4.18 JK–1 g–1. Assume that all of the heat given out is used to heat the water.

5d3 marks

Explain why the enthalpy change for the formation of calcium chloride cannot be measured using mean bond enthalpy data.

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