Calorimetry

A value for enthalpy of solution can be determined from the results of a calorimetry experiment or by the application of Hess’s law.

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

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.

The student calculated the enthalpy of solution of potassium hydroxide

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.

Table 2 shows some enthalpy change data.

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

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

This question is about enthalpy of combustion.

Limonene is found in the skin of citrus fruits.

Figure 6 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).

Table 3 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 Table 3.

This question is about biofuels.

Palmitic acid, CH₃(CH₂)₁₄COOH, can be made by hydrolysis of the triester in palm oil under acidic conditions. Palmitic acid can be used as a biofuel.

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

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

State how the value calculated in Question 08.2 is likely to differ from data book values.

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

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

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

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

Glucose has the structural formula shown.

Table 3 shows some mean bond enthalpy values.

Use the equation and the data in Table 3 to calculate an approximate value of Δ H for the fermentation of glucose. For this calculation you should assume that all the substances are in the gaseous state.

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

The equation for this reaction is

Table 4 shows some standard enthalpies of formation.

Use the data in Table 4 to calculate a standard enthalpy change value for this reaction.

The molar enthalpy of vaporisation (H_vap) of a liquid is the enthalpy change when one mole of liquid is converted to vapour at the boiling point of the liquid.

A student does an experiment to determine H_vap for water.

The student:

• places a large beaker on a balance

• pours 500 cm³ of water into the beaker

• uses a 2.4 kW heater to raise the temperature of the water to 100 C

• records the mass of the beaker and hot water

• uses the 2.4 kW heater to boil the water for 100 s

• records the mass of the beaker and remaining water.

The loss in mass is 103 g

Calculate H_vap for water.

Table 2 shows some data about three compounds that all contain the same number of electrons.

All three compounds in Table 2 are polar. Ethanol is the most polar and ethylamine is the least polar.

Explain why all three molecules are polar and why ethylamine is the least polar. In your answer refer to the shapes around, and relative electronegativities of, the most electronegative atoms.

Explain the trend in the boiling points of the three compounds. Refer to the intermolecular forces in all three compounds in your answer.

A value for enthalpy of solution can be determined in two ways:

• from a cycle, using lattice enthalpy and enthalpies of hydration

• from the results of a calorimetry experiment.

Define the term enthalpy of lattice dissociation.

The enthalpy of solution for ammonium nitrate is the enthalpy change for the reaction shown.

Draw a suitably labelled cycle and use it, with data from Table 1, to calculate the enthalpy of lattice dissociation for ammonium nitrate.

A student does an experiment to determine a value for the enthalpy of solution for ammonium nitrate.

The student uses this method.

• Measure 25.0 cm³ of distilled water in a measuring cylinder.

• Pour the water into a beaker.

• Record the temperature of the water in the beaker.

• Add 4.00 g of solid NH₄NO₃ to the water in the beaker.

• Stir the solution and record the lowest temperature reached.

Table 2 shows the student’s results.

Calculate the enthalpy of solution, in kJ mol¹, for ammonium nitrate in this experiment.

Assume that the specific heat capacity of the solution, c = 4.18 J K¹ g¹

Assume that the density of the solution = 1.00 g cm³

The uncertainty in each of the temperature readings from the thermometer used in this experiment is ±0.1°C

Calculate the percentage uncertainty in the temperature change in this experiment.

Percentage uncertainty

Suggest a change to the student’s method, using the same apparatus, that would reduce the percentage uncertainty in the temperature change.

Give a reason for your answer.

Another student obtained a value of +15 kJ mol¹ using the same method.

Suggest the main reason for the difference between this experimental value for the enthalpy of solution and the correct value of +26 kJ mol¹

Table 3 shows some entropy data at 298 K

	Entropy S / J K1mol1
NH4NO3(s)	151
NH4(aq)	113
NO3(aq)	146

Calculate a value for the Gibbs free-energy change (G), at 298 K, for the reaction when ammonium nitrate dissolves in water.

Use data from Table 3 and the value of H from the equation. Assume for the solvent, water, that the entropy change, S = 0

Explain what the calculated value of G indicates about the feasibility of this reaction at 298 K

Ammonium nitrate decomposes as shown.

The entropy change (S) for this reaction is +144 J K^–1 mol¹

Calculate the temperature at which this reaction becomes feasible.

A student does an experiment to determine a value for the enthalpy of combustion of heptane. Figure 2 shows some of the apparatus used.

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

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.

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

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

A bomb calorimeter can be used for accurate determination of the heat change during combustion of a fuel.

A bomb calorimeter is a container of fixed volume that withstands the change in pressure during the reaction.

The fuel is mixed with pure oxygen in the calorimeter, ignited and the temperature change is recorded.

The total heat capacity (C_cal) of the calorimeter is calculated using a fuel for which the heat change is known.

In an experiment to calculate C_cal, 2.00 g of hexane (M_r = 86.0) is ignited. A temperature change (T) of 12.4 °C is recorded.

Under the conditions of the experiment, 1.00 mol of hexane releases 4154 kJ of energy when combusted.

The heat energy released in the calorimeter, q = C_calT

Calculate the heat capacity (C_cal) in kJ K¹

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 mol¹, for octane in this combustion reaction.

If you were unable to calculate a value for C_cal in Question 05.1, use 6.52 kJ K¹ (this is not the correct value).

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

The thermometer used to measure the temperature change of 12.2 °C in Question 05.2 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.

A student carried out an experiment to find the temperature rise for a reaction between hydrochloric acid and sodium hydroxide solution.

• The student used a measuring cylinder to place 50 cm³ of 0.400 mol dm^–3 hydrochloric acid into a glass beaker.

• The student recorded the temperature at one-minute intervals for three minutes.

• At the fourth minute the student added 50 cm³ of 0.400 mol dm^–3 sodium hydroxide solution and stirred to mix the solutions, but did not record the temperature.

The student recorded the temperature at one-minute intervals for a further eight minutes.

The results are shown in Table 3.

Plot a graph of temperature against time on the grid opposite. Use your graph to find the temperature rise, ∆ T, at the fourth minute. Show your working on the graph by drawing suitable lines of best fit.

The uncertainty in each of the temperature readings from the thermometer used in this experiment was ±0.1° C

Calculate the percentage uncertainty in the value for the temperature rise.

Suggest a change to the experiment that would minimise heat loss.

Suggest and explain another change to the experiment that would decrease the percentage uncertainty in the use of the same thermometer.

A second student completed an experiment to determine the enthalpy of neutralisation for the reaction between ethanedioic acid solution (HOOCCOOH) and potassium hydroxide solution.

The student added 25 cm³ of 0.80 mol dm^–3 ethanedioic acid solution to 75 cm³ of 0.60 mol dm^–3 potassium hydroxide solution. The temperature increased by 3.2 ° C

Give an equation for the reaction between ethanedioic acid solution and potassium hydroxide solution. Calculate the enthalpy change (∆ H) per mole of water formed in this reaction. Assume that the specific heat capacity of the reaction mixture is 4.2 J K^–1g^–1 Assume that the density of the reaction mixture is 1.00 g cm^–3

In a similar experiment to that in Question 04.5, the enthalpy of neutralisation for the reaction between sulfuric acid and potassium hydroxide solution was found to be –57.0 kJ mol^–1 per mole of water formed.

Suggest an explanation for the difference between this value and your answer to Question 04.5. (If you were unable to obtain an answer to Question 04.5 you should assume a value of –28.5 kJ mol^–1. This is not the correct answer.)

Turn over for Section B