The effect of changing conditions on equilibrium

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 equilibrium and rates.

This question is about two oxides of sulfur: sulfur dioxide, SO₂, and sulfur trioxide, SO₃.

SO₃ decomposes to form SO₂ and O₂, as shown in Equilibrium 18.1.

2.25 moles of SO₃ is heated to 550 C in the presence of a catalyst and the resulting mixture allowed to reach equilibrium.

The equilibrium mixture contains 0.900 mol of SO₂ and the total pressure is 2.80 atm.

Calculate the numerical value for K_p for Equilibrium 18.1 under these conditions and state the units of K_p.

Give your answer to 3 significant figures.

The numerical values of K_p for Equilibrium 18.1 at temperatures T₁ and T₂ are shown below.

Explain why T₂ is a higher temperature than T₁.

Suggest how the value of K_p would change if the reaction was repeated with no catalyst added and the pressure of the system increased.

Tick (✓) one box in each row.

SO₂ and SO₃ both have molecules with sulfur in the centre and bond angles of approximately 120.

Explain why the bond angles in SO₃ are 120.

Explain why both SO₂ and SO₃ have polar bonds, but only SO₂ has polar molecules.

This question is about oxides of nitrogen.

An investigation is carried out on the equilibrium system shown below.

A sealed flask containing 6.00 moles of NO₂(g) is heated to a constant temperature and allowed to reach equilibrium.

The equilibrium mixture contains 5.40 mol of NO₂(g), and the total pressure is 5.00 atm.

Determine the value of K_p and give your answer to 3 significant figures.

Include an expression for K_p and the units of K_p in your answer.

The sealed flask in (a)(i) is then heated to a higher temperature at an increased pressure. The system is allowed to reach equilibrium again.

Explain why it is difficult to predict how these changes in reaction conditions affect the amount of N₂O₄(g) formed at equilibrium.

N₂O₄ reacts fully with oxygen to form a different oxide of nitrogen, oxide A, as the only product.

Oxide A is collected and cooled to 75.0 °C at a pressure of 101 kPa.

Under these conditions, oxide A is a gas that occupies a volume of 74.0 cm³ and has a mass of 0.280 g.

Calculate the molar mass of oxide A and suggest its molecular formula.

molar mass =........ g mol

This question is about equilibria involving hydrogen.

* Hydrogen is used industrially to manufacture ammonia.

The equilibrium is shown below.

1.20 mol N₂(g) is mixed with 3.60 mol H₂(g) in a 8.00 dm³ container.

The mixture is heated to 550 °C with an iron catalyst and allowed to reach equilibrium.

The equilibrium mixture contains 0.160 mol of NH₃.

Determine the equilibrium constant K_c for Equilibrium 20.1, and explain why the operational conditions used by industry may be different from those required for a maximum equilibrium

yield of ammonia.

In industry, hydrogen is also used to reduce the iron oxide Fe₃O₄ as shown in Equilibrium 20.2.

The reaction is carried out at 500 °C.

When the temperature is decreased, the value of K_p decreases.

Determine whether the forward reaction is exothermic or endothermic. Explain your answer.

Two students are discussing the effect of pressure on the equilibrium position of Equilibrium 20.2.

Student 1 says:

“There are more moles of products than reactants, so increasing the pressure will shift the equilibrium to the left hand side.”

Student 2 disagrees.

Determine which student is correct. Justify your answer.

Two students plan to investigate Equilibrium 4.1, shown below.

The students are supplied with the equilibrium mixture in Equilibrium 4.1 at room temperature.

• One student heats 20 cm³ of the mixture to 50 °C.

• The other student heats 20 cm³ of the mixture to 90 °C.

The students use colorimetry to observe how the colour of the equilibrium mixture changes over time.

• The colorimeter is set up so that the greater the absorbance, the greater the

The students plot the graph below from the results of the experiment.

Use the graph and relevant chemical theory to answer the following. Include all reasoning:

• Explain the different initial rates at 50 °C and 90 °C.

• Predict the sign of ∆ H for the forward reaction in Equilibrium 4.1.

The students investigate how addition of aqueous silver nitrate, AgNO₃(aq), affects the equilibrium position in Equilibrium 4.1.

The graph shows the changes in the equilibrium concentrations of CoCl₄^2–, Cl^– and [Co(H₂O)₆]²⁺ after addition of the AgNO₃(aq).

The AgNO₃(aq) is added at time = t₁.

Explain why the Cl^– concentration drops sharply at time = t₁.

An equilibrium system is shown below.

Different amounts of CO, NO₂, CO₂ and NO are added to four containers, as shown below.

Which container would have the largest concentration of CO(g) at equilibrium?

Your answer

This question is about chemical equilibrium.

Nitrogen monoxide, NO, and oxygen, O₂, react to form nitrogen dioxide, NO₂, in the reversible reaction shown in Equilibrium 20.1.

A dynamic equilibrium exists in a closed system.

State one other feature of a dynamic equilibrium.

Show that the formation of NO₂ in Equilibrium 20.1 is feasible at 25 °C.

Determine the maximum temperature, in K, for feasibility.

Give your answer to an appropriate number of significant figures.

A chemist investigates the equilibrium shown in Equilibrium 20.1.

The chemist mixes together 1.60 mol of NO(g) and 1.50 mol of O₂(g) in a container and the mixture is allowed to reach equilibrium.

At equilibrium:

• 75% of the NO(g) has been converted to NO₂(g)

• the total pressure is 1.21 MPa.

Calculate K_p, in MPa^–1, for Equilibrium 20.1.

Give your answer to 3 significant figures.

The chemist then repeats the experiment three times. In each experiment, the chemist makes one change but uses the same initial amounts of NO and O₂.

Complete the table to show the predicted effect of each change compared with the original experiment.

Only use the words greater, smaller or same.

State and explain how increasing the temperature affects the position of this equilibrium and whether the forward reaction is exothermic or endothermic.

Plot a graph of ln K_p against using the axes provided on the opposite page. Use your graph and equation 5.1 to determine H, in kJ mol⁻¹, for this equilibrium.

Give your answer to 3 significant figures.

H =........ kJ mol⁻¹

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.

Nitrogen monoxide, NO, and oxygen, O₂, react to form nitrogen dioxide, NO₂, in the reversible reaction shown in equilibrium 18.1.

Write an expression for K_c for this equilibrium and state the units.

K_c =

A chemist mixes together nitrogen and oxygen and pressurises the gases so that their total gas volume is 4.0 dm³.

• The mixture is allowed to reach equilibrium at constant temperature and volume.

• The equilibrium mixture contains 0.40 mol NO and 0.80 mol O₂.

• Under these conditions, the numerical value of K_c is 45.

Calculate the amount, in mol, of NO₂ in the equilibrium mixture.

The values of K_p for equilibrium 18.1 at 298 K and 1000 K are shown below.

Predict, with a reason, whether the forward reaction is exothermic or endothermic.

The chemist increases the pressure of the equilibrium mixture at the same temperature.

State, and explain in terms of K_p, how you would expect the equilibrium position to change.

This question looks at reactions of hydrogen peroxide and of cobalt(II) ions.

Aqueous hydrogen peroxide decomposes as shown in equation 2.1.

The reaction is catalysed by manganese(IV) oxide, MnO₂.

A student investigates the decomposition of a hydrogen peroxide solution as outlined below.

• The student adds 50.00 cm³ of H₂O₂(aq) to a conical flask.

• The student adds a small spatula measure of MnO₂ and quickly connects the flask to a gas syringe.

• The student measures the volume of oxygen every 200 seconds.

The student allows the reaction in equation 2.1 to proceed until no more gas is evolved. The volume of O₂ in the syringe is now 55 cm³, measured at RTP.

Calculate the initial concentration of the H₂O₂.

Give your answer to two significant figures.

initial concentration of H₂O₂ =........ mol dm^–3

Hydrogen peroxide can act as an oxidising agent or as a reducing agent.

Some standard electrode potentials are shown below.

Use this information to write an equation for a reaction in which hydrogen peroxide acts as a reducing agent.

Cobalt(II) forms complex ions with water ligands and with chloride ligands.

• With water ligands, cobalt(II) forms a pink octahedral complex ion, [Co(H₂O)₆]²⁺.

• With chloride ligands, cobalt(II) forms a blue tetrahedral complex ion.

A student dissolves cobalt(II) sulfate in water in a boiling tube. A pink solution forms.

Experiment 1 The student places the boiling tube in a water bath at 100 °C. Concentrated hydrochloric acid is added dropwise. The colour of the solution changes from pink to blue.

Experiment 2 The student places the boiling tube from experiment 1 in an ice/water bath at 0 °C. The colour of the solution changes from blue to pink.

Write the equilibrium equation for the reaction that takes place when the colour of the solution changes.

Explain the observations and predict whether the formation of the blue colour is exothermic or endothermic.