Rates of Reaction: Chemistry 0620 (Cambridge O Level / IGCSE)

A water-testing laboratory checks a sample of tap water for chloride ions by mixing it with silver nitrate solution. A white precipitate of silver chloride forms immediately: $\text{AgNO}_3(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$

A technician wants this precipitation reaction to happen faster.

Which change would not increase the rate of this reaction?

A bakery's quality-control team investigates how quickly yeast produces carbon dioxide gas from sugar in a bread dough mixture: $\text{C}_6\text{H}_{12}\text{O}_6 \xrightarrow{\text{yeast enzymes}} 2\text{C}_2\text{H}_5\text{OH} + 2\text{CO}_2$

They place $150\text{ g}$ of a dough mixture (containing yeast, sugar and water) into an empty flask standing on a mass balance. The empty flask, together with its loosely fitting cotton wool plug, has a mass of $120.00\text{ g}$, so the balance reads $270.00\text{ g}$ as soon as the dough is added. The cotton wool plug lets carbon dioxide gas escape but no dough is lost. The total mass of the flask and its contents is recorded every 2 minutes.

(a) Explain how the readings in the table show that carbon dioxide gas is being produced, and explain why the mass keeps decreasing throughout the experiment. [2]

(b) Calculate the average rate of reaction, in grams per minute, (i) between $t=0$ and $t=2$ minutes, and (ii) between $t=8$ and $t=10$ minutes. [2]

(c) The average rate calculated in (b)(ii) is much smaller than the average rate calculated in (b)(i). State and explain the reason for this, in terms of the sugar in the dough. [1]

(d) The bakery repeats the test with an identical dough mixture, but keeps it in a warmer room instead. Describe one similarity and one difference you would expect between this new mass–time graph and the original one. [2]

A teacher demonstrates the reaction between colourless hydrogen gas and purple iodine vapour, which produces colourless hydrogen iodide gas: $\text{H}_2(g) + \text{I}_2(g) \rightarrow 2\text{HI}(g)$

The gases are mixed inside a sealed gas syringe, and the fading of the purple colour is used to judge how far the reaction has progressed.

(a) The teacher pushes in the plunger of the syringe, reducing the volume of the gas mixture and increasing its pressure, while keeping the temperature the same. Explain, using collision theory, why this increases the rate of reaction. [3]

(b) In a separate trial at the original pressure, the teacher instead increases the temperature of the gas mixture. Explain, using collision theory, why increasing the temperature increases the rate of reaction. Your answer should refer to two separate effects of a higher temperature on the reacting particles. [4]

(c) A finely divided platinum catalyst is then added to the gas mixture, at the original pressure and temperature. The activation energy of the uncatalysed reaction is $170\text{ kJ/mol}$; with the platinum catalyst present, the reaction instead proceeds by a pathway with an activation energy of only $70\text{ kJ/mol}$. Explain why this increases the rate of reaction. [2]

A school science technician prepares nitrogen gas for other experiments by warming a mixture of ammonium chloride solution and sodium nitrite solution, which react together: $\text{NH}_4\text{Cl}(aq) + \text{NaNO}_2(aq) \rightarrow \text{NaCl}(aq) + \text{N}_2(g) + 2\text{H}_2\text{O}(l)$

The technician carries out two trials, each using the same total number of moles of ammonium chloride, and keeping every other condition the same:

• Trial 1: the ammonium chloride is dissolved in a larger volume of water, making a more dilute solution.
• Trial 2: the ammonium chloride is dissolved in a smaller volume of water, making a more concentrated solution.

The volume of nitrogen gas collected in a gas syringe is recorded every 30 seconds.

(a) State the factor that is being investigated by comparing trial 1 and trial 2. [1]

(b) Calculate the average rate of reaction in trial 1 between $t=0$ and $t=60$ seconds. Give the units of your answer. [2]

(c) Both trials eventually produce the same final volume of gas, $32\text{ cm}^3$. Using the fact (given above) that both trials use the same total number of moles of ammonium chloride, state whether it is the rate of reaction or the total amount of gas produced that is changed by altering the concentration of a solution, and explain your answer. [2]

(d) State what the shape of both graphs at $t=150\text{ s}$ tells you about each reaction at this time. [1]

A margarine manufacturer hardens liquid vegetable oil into a solid fat by reacting it with hydrogen gas, using a nickel catalyst: $\text{vegetable oil}(l) + \text{H}_2(g) \xrightarrow{\text{Ni catalyst}} \text{margarine}(s)$

This is an exothermic reaction: the reactants have more energy than the products.

Which statement correctly describes the effect of the nickel catalyst on this reaction?

A student investigates how the size of the pieces of calcium carbonate affects the rate of its reaction with dilute hydrochloric acid: $\text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$

The student carries out two experiments, each using $0.40\text{ g}$ of calcium carbonate and $50\text{ cm}^3$ of the same dilute hydrochloric acid (the acid is in excess in both experiments):

• Experiment 1: the calcium carbonate is in the form of large lumps.
• Experiment 2: the calcium carbonate is in the form of small chips (the same total mass, $0.40\text{ g}$).

The volume of carbon dioxide gas collected in a gas syringe is recorded every $20\text{ s}$.

(a) State the factor being investigated by comparing Experiment 1 and Experiment 2. [1]

(b) Calculate the average rate of reaction in Experiment 2 between $t=0$ and $t=40$ seconds. Give the units of your answer. [2]

(c) Suggest two changes, other than making the pieces of calcium carbonate smaller, that would each increase the rate of reaction in Experiment 1. For each change, name the factor affecting rate that it demonstrates. [2]

(d) State the total volume of gas that Experiment 1 will eventually reach, once its reaction is complete, and explain your answer. [1]

A factory manufactures sulfur trioxide gas by reacting sulfur dioxide gas with oxygen gas over a vanadium(V) oxide catalyst: $2\text{SO}_2(g) + \text{O}_2(g) \rightarrow 2\text{SO}_3(g)$

An engineer wants to increase the rate of this reaction, without changing the catalyst that is used.

Which change would increase the rate of this reaction?

A cleaning-products factory tests whether manganese(IV) oxide can be used as a catalyst to speed up the decomposition of hydrogen peroxide solution into water and oxygen gas: $2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g)$

A technician sets up two trials, each using the same volume and concentration of hydrogen peroxide solution, at the same temperature:

• Trial 1: the hydrogen peroxide solution alone, with no catalyst added.
• Trial 2: the same hydrogen peroxide solution, with a small measured mass of powdered manganese(IV) oxide added at the start.

The volume of oxygen gas collected in a gas syringe is recorded every $20\text{ s}$.

(a) State the factor being investigated by comparing Trial 1 and Trial 2. [1]

(b) Calculate the average rate of reaction in Trial 2 between $t=0$ and $t=20$ seconds. Give the units of your answer. [2]

(c) State what the total volume of oxygen gas collected by the end of each trial will eventually be, once both reactions are complete, and explain your answer. [2]

(d) The technician weighs the manganese(IV) oxide before adding it to Trial 2, and weighs it again (after filtering and drying it) once the reaction has finished. Suggest what this second measurement would show, and explain why this shows that manganese(IV) oxide is acting as a catalyst rather than as a reactant in this reaction. [1]

Two students each investigate a different factor affecting the rate of reaction between magnesium metal and dilute hydrochloric acid, which produces hydrogen gas: $\text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g)$

• Student A compares using $1.0\text{ mol/dm}^3$ hydrochloric acid with $2.0\text{ mol/dm}^3$ hydrochloric acid. The same volume of acid and the same mass and form of magnesium (a single ribbon) are used in both trials, and the acid is in excess in both cases.
• Student B compares using $1.0\text{ g}$ of magnesium as a single ribbon with $1.0\text{ g}$ of magnesium as a powder. The same volume and concentration of hydrochloric acid is used in both trials, and the acid is in excess in both cases.

(a) Explain, using collision theory, why Student A's trial with $2.0\text{ mol/dm}^3$ acid has a greater initial rate of reaction than the trial with $1.0\text{ mol/dm}^3$ acid. Refer to the number of particles per unit volume in your answer. [3]

(b) Explain, using collision theory, why Student B's trial using magnesium powder has a greater initial rate of reaction than the trial using a single ribbon. [3]

(c) For each student's investigation, state whether the factor being changed affects the total volume of hydrogen gas eventually collected, given that the acid is in excess in every trial. [2]

A student heats a reaction mixture from $20\,^{\circ}\text{C}$ ($293\text{ K}$) to $30\,^{\circ}\text{C}$ ($303\text{ K}$). The rate of reaction roughly doubles over this $10\,^{\circ}\text{C}$ rise, even though the average kinetic energy of the particles (which is proportional to absolute temperature) increases by only about $3\%$ over the same range.

Which statement correctly explains why such a small rise in temperature produces such a large increase in rate?