Energetics (AQA A Level Chemistry)

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.

Define the term standard enthalpy of combustion, Δ_cH^ϴ. 

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

Table 1

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.

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

State the meaning of the term mean bond enthalpies. 

Define the term standard enthalpy of formation, Δ_fH^ϴ.

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.

Some enthalpy of formation data are shown below in Table 1. 

Table 1

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. 

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. 

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. 

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 C to 22.3 C. 

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. 

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. 

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. 

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. 

The table below contains some key enthalpy of combustion data. 

Table 1

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

Show all working. 

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)     Δ_rH^ϴ = -789 kJ mol^-1 

The structure of hydrazine is shown in Figure 1. 

Figure 1

Table 2

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

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.

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

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

Table 1

The temperature increased by 34.5 C 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.

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.

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. 

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

Table 2

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

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.

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.

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.

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

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.

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 

Explain why the student’s temperature change of 3.5 C is not an accurate temperature change.

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

Another student completed the same investigation and recorded a maximum temperature of 22.4 C. 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.

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. 

The alkanes are a homologous series of hydrocarbons with the general formula C_nH_2n+2. 

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

Table 1

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

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

Table 2 shows some bond enthalpy data.

Table 2

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

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.)

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

Table 3 shows some standard enthalpy of combustion data.

Table 3

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

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.

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.

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

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.

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.

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.

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

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

Table 1 shows the relevant enthalpy change data.

Table 1

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

The enthalpy change of solution for calcium chloride (M_r 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 C. 

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.

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