Hess's law

This question is about carbon monoxide and carbon as reducing agents in industry.

Iron(III) oxide, Fe₂O₃, can be reduced by carbon monoxide to form iron.

Some standard enthalpy change of formation, _fH, and standard entropy, S, values are shown in the table.

Substance	Fe2O3(s)	CO(g)	Fe(s)	CO2(g)
Δf H / kJ mol–1	–824	–111	0	–394
S / J K–1 mol–1	87	198	To calculate in part (a)(iv)	214

Explain why the standard enthalpy change of formation of Fe has a value of 0 kJ mol^–1.

Calculate the enthalpy change, _rH, for the forward reaction in Equilibrium 17.1.

H =........ kJ mol^–1

Explain why the standard entropy of CO₂ is greater than the standard entropy of Fe₂O₃.

The standard entropy change, S, for the forward reaction in Equilibrium 17.1 has a value of +15 J K^–1 mol^–1.

Use data from the table to calculate the standard entropy, S, of Fe.

S =........ J K^–1 mol^–1

Carbon can also reduce Fe₂O₃(s) to iron as shown in Equilibrium 17.2.

The forward reaction is only feasible at high temperatures.

• Show that the forward reaction is not feasible at 100 °C.

• Calculate the minimum temperature, in K, for the forward reaction to be feasible.

Give the minimum temperature to the nearest whole number.

Steam can be reduced by carbon monoxide in a reversible reaction.

At 700 K the value of K_p is 8.13.

Explain why K_p for Equilibrium 17.3 has no units.

An equal number of moles of steam and carbon monoxide are left at 700 K in a sealed container until equilibrium is reached.

At equilibrium, both H₂ and CO₂ have a partial pressure of 211 kPa.

Calculate the partial pressures of H₂O and CO at equilibrium.

At a second temperature the value of K_p for Equilibrium 17.3 is 129.

Explain whether the second temperature is higher or lower than 700 K.

This question is about energy changes.

Hydrogen peroxide decomposes as shown in Reaction 16.1.

The table shows enthalpy changes of formation and entropies.

	Hf / kJ mol–1	S / J K–1 mol–1
H2O2(l)	–188	110
H2O(l)	–286	70.0
O2(g)	0	205

Calculate the free-energy change, G, in kJ mol^–1, of Reaction 16.1 at 25 C.

Give your answer to 3 significant figures.

G =........ kJ mol^–1

The decomposition of hydrogen peroxide shown in Reaction 16.1 is feasible.

Suggest why Reaction 16.1 does not take place at 25 C despite being feasible.

The rate of decomposition of hydrogen peroxide shown in Reaction 16.1 can be increased by adding a small amount of powdered manganese(IV) oxide, MnO₂.

The MnO₂ acts as a catalyst.

Complete the enthalpy profile diagram for Reaction 16.1 using formulae for the reactants and products.

• Use E_a to label the activation energy without MnO₂.

• Use E_c to label the activation energy with MnO₂.

• Use H to label the enthalpy change of reaction.

Explain why MnO₂ is described as a heterogeneous catalyst for this reaction.

Mn₃O₄ is a compound in which Mn has two different oxidation states. The two oxidation states are different from the Mn in MnO₂.

Suggest the two oxidation states of manganese in Mn₃O₄.

Define the term lattice enthalpy.

Calculate the lattice enthalpy of MnO.

lattice enthalpy =........ kJ mol^–1

This question is about enthalpy changes of reactions involving hydrocarbons.

A student determines the enthalpy change of combustion, _cH, of heptane, C₇H₁₆, using the method outlined below.

• Add 150 g of water to a beaker and measure its temperature.

• Weigh a spirit burner containing heptane and use it to heat the water.

• Extinguish the flame and record the maximum temperature reached by the water.

• Reweigh the spirit burner.

The temperature of the water increased by 10.5 °C.

The spirit burner decreased in mass by 0.133 g.

Use the student’s results to determine the enthalpy change of combustion of heptane, _cH (C₇H₁₆), in kJ mol^–1.

_cH (C₇H₁₆) =........ kJ mol^–1

Nonane, C₉H₂₀, can be broken down by heat to form pentane, C₅H₁₂, and ethene, C₂H₄.

The enthalpy changes of combustion of C₉H₂₀(g) and C₂H₄(g) are shown in the table below.

Use H in Reaction 1 and the enthalpy changes of combustion in the table to determine the enthalpy change of combustion of C₅H₁₂(g).

∆_cH (C₅H₁₂(g)) =........ kJ mol^–1

Ethene can be produced from ethanol, as shown in Reaction 2 below.

Predict the sign of the entropy change, S, for Reaction 2.

Explain your reasoning.

This question is about energy changes.

Magnesium reacts with aqueous silver nitrate, AgNO₃(aq) as shown below.

A student adds an excess of magnesium to 100.0 cm³ of 0.400 mol dm^–3 AgNO₃(aq). The initial temperature is 20.0 °C.

Determine the maximum temperature reached in this reaction. Give your answer to 3 significant figures.

Assume that the specific heat capacity and density of the solution are the same as for water, and that there are no heat losses.

The student wants to repeat the experiment, but there is not enough AgNO₃(aq) left to use another 100.0 cm³ portion.

The student decides to modify the method by adding an excess of magnesium to 50.0 cm³ of 0.400 mol dm^–3 AgNO₃(aq).

Predict, with reasons, how this modification would affect the maximum temperature reached. Assume that there are no heat losses.

Nitric acid is manufactured from ammonia in a multi-stage process. The equation for the first stage in this process is shown in Reaction 17.1.

Some standard enthalpy changes of formation are shown in the table.

Compound	f H / kJ mol–1
NH3(g)	–46
H2O(l)	–286

Explain the term enthalpy change of formation.

Calculate the standard enthalpy change of formation, _fH, of NO(g).

_fH of NO(g) =........ kJ mol^–1

Carbon disulfide, CS₂, reacts with dinitrogen oxide, N₂O, as shown in Reaction 17.2.

Standard entropies, S, are shown in the table.

Substance	CS2(l)	N2O(g)	S8(s)	CO2(g)	N2(g)
S/ J K–1 mol–1	151	220	256	214	192

Explain the term entropy.

The free energy change, G, of Reaction 17.2 is –2672 kJ mol^–1 at 25 °C.

Calculate the enthalpy change, H, of Reaction 17.2, in kJ mol^–1.

H =........ kJ mol^–1

A student concludes that Reaction 17.2 is feasible at all temperatures.

Explain whether the student is correct or not.

Propane, C₃H₈, (boiling point −42 °C) is used as ‘camping gas’. A student plans to determine the enthalpy change of combustion of propane, _cH (C₃H₈), by two methods.

The student first carries out an experiment using the apparatus below.

Determine the enthalpy change of combustion of propane, in kJ mol⁻¹.

Give your answer to 3 significant figures.

_cH (C₃H₈) =........ kJ mol⁻¹

The student finds that the experimental enthalpy change _cH (C₃H₈) is much less exothermic than the accurate standard enthalpy change _cH (C₃H₈) in databases.

One reason could be that the student’s experiment had not been carried out under standard conditions.

Suggest two other reasons for this difference in enthalpy change.

Sulfuric acid is an important chemical used to make detergents, fertilisers and dyes. It is manufactured in a multi-step process.

In the first step of the manufacture of sulfuric acid, sulfur dioxide, SO₂, can be made from the combustion of hydrogen sulfide, H₂S, shown in Reaction 1.

Explain why the enthalpy change for Reaction 1 has a negative value.

Use ideas about enthalpy changes associated with bond breaking and bond making.

Some standard entropy values are given below.

Substance	H2S(g)	O2(g)	SO2(g)	H2O(l)
S / J K−1 mol−1	206	205	248	70

Using calculations, explain whether Reaction 1 is feasible at 20 °C.

Calculations

Explanation for feasible or non feasible........

Write an expression for K_p for Equilibrium 1.

Include the units.

Determine the value of K_p for Equilibrium 1 at 1000 K.

Show all your working.

Give your answer to 3 significant figures.

The chemist repeats the experiment in (b) at a different temperature.

The chemist finds that the value of K_p is greater than the answer to (b)(ii).

Explain whether the temperature in the second experiment is higher or lower than 1000 K.

Vanadium(V) oxide, V₂O₅(s), is used as a catalyst in equilibrium 1.

Explain whether vanadium(V) oxide is acting as a homogeneous or heterogeneous catalyst.

This question is about copper(II) sulfate, CuSO₄, and sodium thiosulfate, Na₂S₂O₃.

* Calculate ∆_solH(CuSO₄(s)) for reaction 5.2 and determine the enthalpy change of reaction 5.1, ∆_rH.

Assume that the specific heat capacity, c, of the solution is the same as for water.

Show your working, including an energy cycle linking the enthalpy changes.

The standard enthalpy change of reaction, ∆_r , and the standard free energy change, ∆ for converting anhydrous sodium thiosulfate to hydrated sodium thiosulfate are shown below.

Standard entropies are given in the table.

Determine the standard entropy, , of anhydrous sodium thiosulfate, Na₂S₂O₃(s).

Give your answer to 3 significant figures.

Sodium thiosulfate contains the thiosulfate ion, S₂O₃²⁻. The displayed formula of S₂O₃²⁻ can be shown as below.

Predict the O–S–S bond angle and name of the shape of the thiosulfate ion.

Bond angle........

Name of shape........

In some of its reactions, the thiosulfate ion forms the tetrathionate ion, S₄O₆²⁻.

The S₄O₆²⁻ ion is a ‘dimer’ of S₂O₃²⁻.

Draw a displayed formula for the S₄O₆²⁻ ion.