Reversible Reactions and Equilibrium: Chemistry 0620 (Cambridge O Level / IGCSE)

Pink hydrated cobalt(II) chloride crystals, $\text{CoCl}_2 \cdot 6\text{H}_2\text{O}$, turn blue when heated strongly, as water is driven off. The blue solid turns pink again when water is added back to it.

Which equation correctly represents this reversible change?

Blue hydrated copper(II) sulfate crystals, $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$, contain water of crystallisation.

A 25.0 g sample of these hydrated crystals contains 9.00 g of water of crystallisation; the remaining 16.0 g is anhydrous copper(II) sulfate.

(a) A second sample of the same hydrated crystals has a mass of 60.0 g. Assuming this sample has the same composition as the first, calculate the mass of water of crystallisation it contains, and hence the mass of anhydrous copper(II) sulfate that would remain if it were heated to constant mass. [2]

(b) State the colour change observed as the 60.0 g sample is heated strongly until no further change occurs, and name the white solid formed. [2]

(c) Write a balanced symbol equation, including state symbols and the $\rightleftharpoons$ symbol, to represent this reversible reaction. [2]

(d) A few drops of water are then added to the white solid formed in (b). Describe what would be observed, and state what this shows about the reaction. [2]

A chemical company is testing a new catalysed reaction between two gaseous elements, $\text{X}_2$ and $\text{Y}_2$, which combine to form the industrially useful compound $\text{XY}_3$:

$\text{X}_2(g) + 3\text{Y}_2(g) \rightleftharpoons 2\text{XY}_3(g)$

The table shows the percentage yield of $\text{XY}_3$ at equilibrium, measured in laboratory-scale trials at different temperatures and pressures.

(a) Use the data in the table to deduce whether the forward reaction is exothermic or endothermic. Explain your reasoning. [3]

(b) Use the data in the table to explain, in terms of the number of gas molecules on each side of the equation, why increasing the pressure increases the percentage yield of $\text{XY}_3$. [3]

(c) The reaction is very slow at 300 °C without a catalyst. State and explain the effect of adding a suitable catalyst on (i) the rate at which equilibrium is reached, and (ii) the percentage yield of $\text{XY}_3$ shown in the table. [2]

(d) Suggest one economic reason, based on the data in the table, why the company would choose to operate at 300 °C rather than 500 °C, given that both temperatures use the same pressure of 200 atm. [2]

A general reversible reaction reaches dynamic equilibrium in a sealed container at constant temperature:

$\text{P}(g) + \text{Q}(g) \rightleftharpoons \text{R}(g) + \text{S}(g)$

Which statement about the equilibrium mixture is correct?

Ammonia is manufactured industrially by the Haber process, using the typical conditions of 450 °C, 200 atmospheres (20 000 kPa) and an iron catalyst:

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$

The forward reaction is exothermic.

A newly-appointed plant manager wants to reduce operating costs and makes two suggestions.

Suggestion 1: lower the pressure from 200 atmospheres to 50 atmospheres, since high-pressure equipment is expensive to build and maintain.

Suggestion 2: raise the temperature from 450 °C to 650 °C, since a hotter reaction mixture reacts faster, so ammonia could be produced more quickly.

(a) State whether Suggestion 1 would give a good yield of ammonia, and explain your answer by referring to the effect of pressure on the position of equilibrium. [3]

(b) State whether Suggestion 2 is a good idea, and explain your answer by referring to the effect of temperature on both the rate of reaction and the equilibrium yield of ammonia. [3]

(c) State one reason, other than yield, why 450 °C is chosen as a compromise rather than the manager's proposed 650 °C. [1]

A small amount of white ammonium chloride solid is placed in the closed end of a long test tube. The tube is held at a slight angle and heated strongly at the closed end, while the open end of the tube remains much cooler.

(a) Describe what would be observed at the heated end of the tube, and further along the tube in the cooler region, as heating continues. [2]

(b) Ammonium chloride decomposes on heating to form two colourless gases, ammonia and hydrogen chloride. Write a balanced symbol equation, including state symbols and the $\rightleftharpoons$ symbol, for this reaction. [2]

(c) Explain, in terms of the forward and reverse reactions taking place in different parts of the tube, why this experiment provides evidence that the decomposition of ammonium chloride is reversible. [2]

Four students each describe an observation from a school experiment.

Student 1: heats green copper(II) carbonate strongly; it turns black and releases a colourless gas that turns limewater cloudy.

Student 2: burns a short piece of magnesium ribbon in air; a white powder is produced and a bright white light is given out.

Student 3: heats blue hydrated copper(II) sulfate crystals, $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$; the crystals turn white, and colourless liquid condenses on the cooler part of the test tube. Adding a few drops of water to the white solid turns it blue again.

Student 4: adds dilute hydrochloric acid to solid sodium carbonate; bubbles of gas are produced and the solid dissolves completely.

Which student's observation is evidence of a reversible reaction?

An unfamiliar reversible reaction between two gases, $\text{D}$ and $\text{E}$, reaches dynamic equilibrium in a closed container at constant temperature:

$2\text{D}(g) + \text{E}(g) \rightleftharpoons 2\text{F}(g)$

(a) A small additional amount of gas $\text{E}$ is injected into the container, while the temperature and the volume of the container are kept constant. Predict, using Le Chatelier's principle, the effect this has on the position of equilibrium, and state what happens to the concentration of $\text{D}$ as a new equilibrium is reached. [3]

(b) $\text{F}(g)$ is then removed continuously from the container as it forms, by a separate process that does not affect $\text{D}$ or $\text{E}$, while the temperature is kept constant. Predict and explain the effect of continuously removing $\text{F}(g)$ on the overall amount of $\text{F}$ produced over time, compared with simply allowing the mixture to settle at a single, static equilibrium. [2]

(c) Starting again from the original equilibrium mixture, the volume of the container is instead doubled at constant temperature, without adding or removing any gas. Using the number of moles of gas on each side of the equation, predict and explain the effect this has on the position of equilibrium. [2]

The Contact process manufactures sulfur trioxide from sulfur dioxide and oxygen, using a vanadium(V) oxide catalyst:

$2\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\text{SO}_3(g)$

The forward reaction is exothermic. Industrially, the Contact process operates at a temperature of 450 °C and a pressure of only about 200 kPa (roughly 2 atmospheres). Much lower than the 20 000 kPa used in the Haber process for ammonia,

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$

even though both reactions have fewer gas molecules on the product side than on the reactant side.

(a) For each reaction, state the number of moles of gas molecules on the reactant side and on the product side, and hence the overall decrease in the number of moles of gas molecules as the forward reaction proceeds. [2]

(b) Use your answer to part (a) to suggest why increasing the pressure has a smaller effect on the position of equilibrium, and therefore on the percentage yield, in the Contact process than in the Haber process. [2]

(c) The Contact process already achieves a very high percentage yield of sulfur trioxide at only 200 kPa. Using your answer to part (b), explain why it would not be economically sensible to increase the pressure further. [2]

(d) State and explain the effect that raising the temperature above 450 °C would have on the equilibrium yield of sulfur trioxide in the Contact process. [2]

The following reaction reaches dynamic equilibrium in a sealed container at constant temperature and constant volume:

$2\text{L}(g) + \text{M}(g) \rightleftharpoons 2\text{N}(g)$

The forward reaction is exothermic.

A small additional amount of gas $\text{L}$ is then injected into the container, while the temperature and the volume of the container are kept constant.

Which statement correctly describes the effect on the position of equilibrium?