Entropy and Gibbs free energy

Nitrogen dioxide gas is formed when magnesium nitrate is heated.

Table 4 shows some thermodynamic data.

Give the standard conditions indicated by the symbol in Δ_f and

Use the equation and the data in Table 4 to calculate the temperature at which this reaction becomes feasible.

Nitrogen dioxide can also be produced in a reaction between ammonia and oxygen.

Table 5 shows some bond enthalpy values.

Assume that the bonds between N and O in NO₂ are the same.

Calculate the mean bond enthalpy, in kJ mol^–1, for a bond between N and O in NO₂

An alloy of platinum and rhodium is used as a catalyst in the reaction.

What type of catalyst is this?

Some ammonia and oxygen are reacted and the mixture is left to reach equilibrium at a fixed temperature.

Explain why increasing the temperature of this equilibrium mixture decreases the mole fraction of NO₂ in the equilibrium mixture.

This question is about enthalpy changes.

Theoretical values for enthalpies of lattice dissociation can be calculated using a perfect ionic model.

State the meaning of the term perfect ionic model.

Enthalpies of lattice dissociation can also be obtained from Born–Haber cycles.

Figure 3 shows an incomplete Born–Haber cycle for the formation of sodium oxide.

Complete Figure 3 by writing formulas, including state symbols, of the appropriate species on each of the two blank lines.

Table 4 shows some enthalpy changes.

Use the data in Table 4 to calculate the enthalpy of lattice dissociation of sodium oxide.

Explain why the second electron affinity of oxygen has a positive value.

Explain why the enthalpy of lattice dissociation for sodium oxide is greater than the enthalpy of lattice dissociation for sodium chloride.

Sodium chloride dissolves in water.

Table 5 shows some more enthalpy changes.

Use the data in Table 5 to calculate the enthalpy of solution for sodium chloride.

Give a reason why data books do not contain a value for the enthalpy of solution of sodium oxide.

Calculate the temperature, in °C, at which this reaction becomes feasible.

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.

This question is about thermodynamics. Consider the reaction shown.

Table 7 shows some thermodynamic data.

Explain why the standard entropy value for carbon dioxide is greater than that for carbon.

State the temperature at which the standard entropy of aluminium is 0 J K^–1 mol^–1

Use the equation and the data in Table 7 to calculate the minimum temperature, in K, at which this reaction becomes feasible.

Methanol is formed when carbon dioxide and hydrogen react.

Table 5 contains enthalpy of formation and entropy data for these substances.

Use the equation and the data in Table 5 to calculate the Gibbs free-energy change (ΔG), in kJ mol^–1, for this reaction at 890 K

Figure 4 shows how the Gibbs free-energy change varies with temperature in a different gas phase reaction.

The straight line graph for this gas phase reaction has been extrapolated to zero Kelvin.

Use the values of the intercept and gradient from the graph in Figure 4 to calculate the enthalpy change (ΔH), in kJ mol^–1, and the entropy change (ΔS), in J K^–1 mol^–1, for this reaction.

State what Figure 4 shows about the feasibility of the reaction.

Figure 1 shows an incomplete Born–Haber cycle for the formation of caesium iodide. The diagram is not to scale.

Figure 1

Table 1 gives values of some standard enthalpy changes.

Complete Figure 1 by writing the formulas, including state symbols, of the appropriate species on each of the two blank lines.

Use Figure 1 and the data in Table 1 to calculate the standard enthalpy of atomisation of iodine.

The enthalpy of lattice formation for caesium iodide in Table 1 is a value obtained by experiment. The value obtained by calculation using the perfect ionic model is –582 kJmol^–1

Deduce what these values indicate about the bonding in caesium iodide.

Use data from Table 2 to show that this reaction is not feasible at 298 K

The equation for the reaction between ammonia and oxygen is shown.

Some standard entropies are given in Table 3.

Calculate the entropy change for the reaction between ammonia and oxygen.

Calculate a value for the Gibbs free-energy change (ΔG), in kJ mol^–1, for the reaction between ammonia and oxygen at 600 °C

(If you were unable to obtain an answer to Question 03.1, you should assume that the entropy change is 211 J K^–1 mol^–1. This is not the correct answer.)

The reaction between ammonia and oxygen was carried out at a higher temperature.

Explain how this change affects the value of Δ G for the reaction.

Platinum acts as a heterogeneous catalyst in the reaction between ammonia and oxygen. It provides an alternative reaction route with a lower activation energy.

Deduce the change in oxidation state of nitrogen, when NH₃ is oxidised to NO

When ammonia reacts with oxygen, nitrous oxide (N₂O) can be produced instead of NO

Give an equation for this reaction.