Moles and amount of substance

A student uses this apparatus to investigate the energy content of different fuels.

This is the student’s method.

• pour some water into the copper can

• record the mass of the spirit burner and fuel

• measure the initial temperature of the water

• place the spirit burner under the copper can and light the burner

• stop heating the water when the temperature reaches 30C

• record the new mass of the spirit burner and fuel

The student repeats the experiment with different fuels.

Explain two variables the student should control to make this a valid test.

In one of the experiments, the student uses ethanol as the fuel.

These are the student’s results for ethanol.

Calculate the value of heat energy change (Q) in joules.

[for water, c = 4.2 J/g/C 1.0 cm³ of water has mass = 1.0 g]

Calculate the enthalpy change (Δ H) in kJ/mol.

[for ethanol, M_r = 46]

Include a sign in your answer.

ΔH =_........ kJ/mol

A student uses this apparatus to investigate the heat energy released when a liquid fuel burns.

This is the student’s method.

• measure the mass of the spirit burner and fuel

• add 100 g of water to the aluminium can

• measure the temperature of the water

• use the spirit burner to heat the water until the temperature rises by 50 °C

• measure the new mass of the spirit burner and fuel

Give a reason why the student uses an aluminium can rather than a glass beaker.

When the fuel burns, a black solid forms on the aluminium can.

Identify the black solid.

Give a reason why the black solid forms.

Show that the heat energy needed to raise the temperature of 100 g of water by 50 °C is approximately 20 000 J.

[for water, c = 4.2 J/g/°C]

The student burns 1.84 g of ethanol.

Calculate the molar enthalpy change, Δ H, in kJ/mol, for the combustion of ethanol.

Include a sign in your answer.

[for ethanol, M_r = 46]

ΔH =_........ kJ/mol

In a different experiment, the student burns 3.7 g of butanol.

This is an equation for the incomplete combustion of butanol.

Complete the equation for the incomplete combustion of butanol.

Calculate the total amount, in moles, of gas produced by the incomplete combustion of 3.7 g of butanol.

This question is about alcohols and carboxylic acids.

These are the structural formulae of alcohol A and carboxylic acid B.

Alcohol A
CH₃CH₂CH₂OH

Carboxylic acid B CH₃(CH₂)_nCOOH

Name alcohol A.

Draw the displayed formula of alcohol A.

Carboxylic acid B has a chain of carbon atoms with no branches.

The number of CH₂ units is represented by the letter n.

Calculate the value of n.

[for carboxylic acid B, M_r = 242]

Alcohol A and carboxylic acid B react together to form an ester.

Give the other product of the reaction.

These are the structural formulae of dicarboxylic acid C and diol D.

Dicarboxylic acid C HOOCCH₂CH₂COOH

Diol D
HOCH₂CH₂OH

This is the equation for the reaction between dicarboxylic acid C and sodium hydroxide solution.

HOOCCH₂CH₂COOH + 2NaOH NaOOCH₂CH₂COONa + 2H₂O

25.0 cm³ of 0.150 mol/dm³ sodium hydroxide solution is completely neutralised by 17.5 cm³ of a solution of dicarboxylic acid C.

Calculate the concentration, in mol/dm³, of the solution of dicarboxylic acid C.

Give your answer to three significant figures.

concentration =_........ mol/dm

Dicarboxylic acid C and diol D react to form a polyester.

Draw the displayed formula of the repeat unit of this polyester.

Diamond and graphite are both forms of the element carbon.

Diamond and graphite both have covalent bonds and giant covalent structures.

The diagram represents the structure of diamond and the structure of graphite.

Give a reason why diamond is an element.

Describe the forces of attraction in a covalent bond.

Explain why graphite conducts electricity.

Explain why diamond is hard but graphite is soft.

Another form of carbon has molecules with the formula C_x

x represents the number of carbon atoms in each molecule.

Each molecule of C_x has a mass of 1.40 × 10⁻²¹ g.

One mole of C_x contains 6.02 × 10²³ molecules.

Calculate the M_r of C_x and the value of x

[for carbon, A_r = 12]

This question is about the metal tantalum, Ta.

Tantalum metal can be produced by heating tantalum chloride (TaCl₅) and hydrogen gas in a furnace.

The other product of the reaction is hydrogen chloride.

Complete the equation for the reaction.

As tantalum chloride is heated, the mass of solid in the furnace decreases leaving tantalum as the only solid product.

The table shows the mass of solid in the furnace at one-hour intervals.

State how the data in the table shows that the reaction is complete.

Use the data to show that the formula of tantalum chloride is TaCl₅

[for tantalum, A_r = 181 for chlorine, A_r = 35.5]

QUESTION 12 CONTINUES ON NEXT PAGE.

Another method of extracting tantalum is by reacting tantalum(V) oxide with carbon.

This is the equation for the reaction.

Explain why this is a redox reaction.

2000 mol of tantalum(V) oxide is heated with 500 000 g of carbon.

Show by calculation that the carbon is in excess.

[for carbon, A_r = 12]

Calculate the maximum mass, in grams, of tantalum that can be obtained from 2000 mol of tantalum(V) oxide.

[for tantalum, A_r = 181]

A scientist reacts lead(II) nitrate solution with sodium bromide solution.

This is the equation for the reaction.

Pb(NO₃)₂(aq) + 2NaBr(aq) PbBr₂(s) + 2NaNO₃(aq)

Describe how the scientist could obtain a pure dry sample of lead(II) bromide (PbBr₂) from the reaction mixture.

The scientist reacts an excess of lead(II) nitrate solution with 25 cm³ of sodium bromide solution of concentration 2.0 mol/dm³.

Show that the amount of sodium bromide used is 0.050 mol.

Show that the maximum theoretical mass of lead(II) bromide is approximately 9 g.

[for PbBr₂ M_r = 367]

The scientist electrolyses molten lead(II) bromide using this apparatus.

Explain why lead(II) bromide needs to be molten rather than solid for electrolysis to occur.

The electrolyte is at a temperature of 400C.

Explain a suitable material for the electrodes.

Give the half‑equation that occurs at the negative electrode.

Bromine forms at the positive electrode.

This is the half‑equation for the reaction at the positive electrode.

State what is observed at the positive electrode.

State why the half‑equation represents an oxidation reaction.

Diamond and graphite are made of carbon atoms, joined together by covalent bonds.

The diagram shows their structures.

State, in terms of electrostatic attractions, what is meant by a covalent bond.

Explain why diamond has a high melting point.

Explain why graphite is a good conductor of electricity.

C₆₀ fullerene is a molecule made of 60 carbon atoms.

The diagram shows the structure of C₆₀ fullerene.

One mole of atoms contains 6.010²³ atoms.

Determine the number of atoms in one mole of C₆₀ fullerene.

Give your answer in standard form.

A scientist uses this apparatus in an experiment to reduce a metal oxide to a metal.

Before heating the mass of the empty tube and the mass of the tube and the metal oxide are recorded.

After heating, the tube is allowed to cool and the mass of the tube and its contents is recorded again.

State why the reaction of the metal oxide to form a metal is described as a reduction reaction.

State why it is important to relight the hydrogen at the end of the tube if the flame goes out.

Explain why it is important to continue passing hydrogen into the tube and burning the hydrogen at the end of the tube until the contents have cooled.

Describe what should be done next to ensure that all the metal oxide has been converted into the metal.

In this experiment a mass of 4.14 g of metal is formed from 4.46 g of the metal oxide.

Calculate the amount, in moles, of oxygen atoms in the sample of the metal oxide.

The formula of the metal oxide is MO, where M represents the symbol of the metal.

Deduce the amount, in moles, of M in the sample of the metal oxide.

Calculate the relative atomic mass of M.

Use the Periodic Table to identify metal M.

A teacher uses this apparatus to test the products of the combustion of liquid hydrocarbons.

Explain the change in appearance that occurs in the limewater.

The equation represents the complete combustion of an alkene.

C_nH_2n + oxygen carbon dioxide + nH₂O

Complete combustion of 0.0100 mol of the alkene produces 2.16 g of water.

Determine the molecular formula of this alkene.

[for H₂O, M_r = 18]

Give a reason why the mass of pure water that collects in the U-tube is less than 2.16 g.

Give a physical test to show that the water that collects in the U-tube is pure.

The teacher burns 30.0 g of heptane.

This is the equation for the complete combustion of heptane.

Calculate the minimum mass of oxygen needed for the complete combustion of 30.0 g of heptane.

[for C₇H₁₆, M_r = 100 for O₂, M_r = 32]

This question is about different hydrated forms of sodium sulfate.

A compound has the formula Na₂SO₄.7H₂O

How many different elements are there in the formula Na₂SO₄.7H₂O?

What is the total number of atoms in the formula Na₂SO₄.7H₂O?

Another hydrated form of sodium sulfate has the formula Na₂SO₄.xH₂O

A student uses this apparatus to find the value of x.

This is the student’s method.

• find the mass of an empty tube

• add solid hydrated sodium sulfate to the tube

• find the mass of the tube and hydrated sodium sulfate

• heat the tube for several minutes

• allow the tube to cool and find the mass of the tube and contents

Describe what the student should do next to make sure that all the water is removed from the hydrated sodium sulfate.

Explain the role of the ice in the beaker.

Describe how the student could prove that the liquid collected is pure water.

The table gives the student’s results.

Use the student’s results to calculate the value of x.

[for Na₂SO₄, M_r = 142 for H₂O, M_r = 18]

Zinc reacts with dilute hydrochloric acid to form zinc chloride and hydrogen gas.

This is the equation for the reaction.

In an experiment, 20 cm³ of hydrochloric acid containing 0.0036 mol are reacted with 1.3 g of zinc granules at a temperature of 30 °C.

Show by calculation that the zinc is in excess.

The volume of hydrogen collected is measured at regular time intervals.

Curve A shows the results of this experiment.

The experiment is repeated using 1.3 g of zinc powder instead of zinc granules.

All other conditions are kept the same.

On the grid, draw the curve you would expect to obtain.

In the original experiment, 20 cm³ of hydrochloric acid containing 0.0036 mol were reacted with 1.3 g of zinc granules at a temperature of 30 °C and curve A was obtained.

The student does two more experiments and obtains curves B and C.

In one of these experiments the student repeats the original method but at a temperature of 20 °C.

Explain in terms of particle collision theory why the curve obtained could be curve B.

In the other experiment the student repeats the original method but uses 20 cm³ of hydrochloric acid containing 0.0018 mol.

Explain why curve C shows the results the student obtained.

Catalysts can be used to speed up reactions.

Describe how a catalyst works.

A student uses this method to investigate the temperature change when solid sodium hydrogencarbonate is added to ethanoic acid solution.

• pour 100 cm³ of ethanoic acid solution into a polystyrene cup

• record the temperature of the ethanoic acid solution

• add a 1 g portion of sodium hydrogencarbonate to the ethanoic acid solution and stir

• record the new temperature

• add further 1 g portions of sodium hydrogencarbonate, stirring and recording the temperature after each portion is added

The table shows the student’s results.

Plot the student’s results on the grid.

Draw a circle around the anomalous result.

Draw a curve of best fit.

Explain why using a polystyrene cup makes the results more accurate.

Suggest a mistake the student might have made to cause the anomalous result.

State how the results show the reaction is complete.

State how the results show that the reaction is endothermic.

Use the results to calculate the heat energy change, Q, in joules.

[for 1.0 cm³ of ethanoic acid solution, mass = 1.0 g]

[for ethanoic acid solution, c = 4.2 J / g / °C]

The student repeats the experiment starting with a different volume of ethanoic acid solution.

The student uses 7.0 g of sodium hydrogencarbonate to neutralise the ethanoic acid solution.

The heat energy change, Q, is 3200 J.

Calculate the molar enthalpy change, Δ H, in kJ / mol.

Include a sign with your answer.

[for sodium hydrogencarbonate, M_r = 84]

ΔH =_........ kJ / mol

Methane reacts with steam to form carbon monoxide and hydrogen.

This is the equation for the reaction.

State why carbon monoxide is poisonous to humans.

Explain the effect, if any, on the yield of hydrogen at equilibrium when a nickel catalyst is used.

The reaction conditions for this reaction are a temperature of 700 °C and a pressure of 5 atmospheres.

The temperature of the reaction mixture is reduced to 600 °C, but the pressure is kept at 5 atmospheres.

Explain the effect on the yield of hydrogen at equilibrium.

The pressure of the reaction mixture is reduced to 4 atmospheres, but the temperature is kept at 700 °C.

Explain the effect on the yield of hydrogen at equilibrium.

Calculate the volume, in dm³, of methane gas at rtp needed to produce 6.6 tonnes of hydrogen gas.

[at rtp, molar volume = 24 dm³ 1 tonne = 10⁶g]

Give your answer in standard form.

This question is about hydrated compounds.

Crystals of hydrated barium chloride (BaCl₂. xH₂O) contain water of crystallisation.

A student uses this apparatus to remove and collect the water from some crystals.

This is the student’s method.

Step 1 record mass of tube A when empty

Step 2 place a sample of hydrated barium chloride crystals in tube A and record new mass

Step 3 heat tube A

Step 4 allow tube A to cool and record mass

Repeat steps 3 and 4 until the mass recorded in step 4 is constant.

These are the student’s results.

mass of tube A = 10.55 g

mass of tube A and BaCl₂. xH₂O = 16.65 g

final mass of tube A and BaCl₂ = 15.75 g

Give a reason why the student repeats steps 3 and 4 until the mass is constant.

Calculate the mass of BaCl₂ that forms in tube A.

Calculate the mass of water lost.

Determine the value of x in BaCl₂. xH₂O

Show your working.

[for BaCl₂, M_r = 208 for H₂O, M_r = 18]

Describe a physical test to show that the water in tube B is pure.

A sample containing 0.02 mol of hydrated copper(II) sulfate is heated using the same apparatus.

The products of the reaction are anhydrous copper(II) sulfate and water.

This is the equation for the reaction.

Give the meaning of the symbol

Describe how the reaction can be used to show that a liquid contains water.

Calculate the maximum number of water molecules in tube B after the sample of hydrated copper(II) sulfate has completely reacted.

One mole of any substance contains 610²³ particles.

When 10.0 kg of a solid fuel is burned, the heat energy produced is 3.28 × 10⁵kJ.

Calculate the heat energy, in kJ/mol, produced by the fuel.

You should assume that the fuel is pure carbon.

Give your answer to three significant figures.

[for carbon, A_r = 12]

heat energy =_........ kJ/mol

The fuel actually contains some impurities, including sulfur.

There is 600 g of sulfur in 20 kg of fuel.

Calculate the percentage of sulfur in the fuel.

When the fuel burns, the 600 g of sulfur produces sulfur dioxide gas.

Calculate the volume, in cm³, of sulfur dioxide gas that would be produced at room temperature and pressure (rtp) when the fuel burns.

[for a gas, molar volume = 24 dm³ at rtp]

State the name of the environmental problem caused by sulfur dioxide in the atmosphere.

The diagram shows the positions of some elements in part of the Periodic Table.

Na A l S C l

K Xe

Give the symbol of a metal from the diagram.

Give the symbol of an element from the diagram that forms an acidic oxide.

Give a similarity in the electron configurations of A l and I n.

Explain which element in the diagram is unreactive.

A teacher adds a small piece of sodium to a glass trough containing water and universal indicator.

The universal indicator changes colour.

The equation for the reaction is

Explain the final colour of the universal indicator.

The teacher repeats the experiment with potassium instead of sodium.

Give one similarity and one difference observed with potassium.

similarity

difference

The reaction with sodium produces 0.036 g of hydrogen gas.

One mole of hydrogen gas contains 6.010²³ molecules.

Calculate the number of molecules of hydrogen gas produced in the reaction with sodium.

Give your answer to two significant figures.

This question is about metal carbonates.

When heated, some metal carbonates decompose to form a metal oxide and carbon dioxide gas.

A student is given three solid metal carbonates, a timer, some limewater and this apparatus.

Describe a method the student can use to find out which metal carbonate decomposes fastest when heated.

A student is given a solid metal carbonate with the formula XCO₃

X represents the symbol of a Group 2 metal.

A student uses this apparatus to heat a sample of XCO₃ until it all decomposes.

The equation for the decomposition of XCO₃ is

The student records the mass of XCO₃ and the mass of carbon dioxide that escapes through the cotton wool plug.

These are the student’s results.

Give a reason why the student uses a cotton wool plug.

Calculate the amount, in mol, of carbon dioxide produced.

[for carbon dioxide M_r = 44]

Use the equation to determine the amount, in mol, of XCO₃ that decomposed.

Use the mass of XCO₃ and your answer to (b)(iii) to calculate the relative formula mass (M_r) of XCO₃

Use your answer to (b)(iv) and the Periodic Table on page 2 to determine the identity of the Group 2 metal X.

Show your working.

This question is about ammonia and ammonium compounds.

Ammonia (NH₃) is a simple covalent molecule.

Draw a dot-and-cross diagram to show the bonding in a molecule of ammonia.

The table shows the names and formulae of some ammonium compounds.

Complete the table by giving the missing information.

When ammonia reacts with sulfuric acid, ammonium sulfate is formed.

Write a chemical equation for this reaction.

Describe a test for ammonium ions.

The table gives some information about ammonia and ammonium compounds.

Calculate the percentage of nitrogen in ammonium nitrate.

[M_r of NH₄NO₃ = 80]

Fertilisers add nitrogen to the soil to help plants grow.

Ammonia and ammonium sulfate can both be used as fertilisers.

Discuss the advantages and disadvantages of using each of these compounds as fertilisers.

Use information from the table in your answer.

[pH of rainwater is approximately 5.6]

This question is about alkanes.

Which of these is the molecular formula of an alkane?

Which of these has the same empirical formula and molecular formula?

In the presence of ultraviolet radiation, methane reacts with bromine to form bromomethane and hydrogen bromide.

State the name of this type of reaction.

Give a chemical equation for this reaction.

One mole of an alkane burns completely in oxygen.

The equation represents the reaction.

The numbers x, y and z are used to balance the equation.

The complete combustion of one mole of the alkane produces 220 g of carbon dioxide and 108 g of water.

Calculate the values of y and z.

[M_r of CO₂ = 44 M_r of H₂O = 18]

Determine the molecular formula of the alkane and the value of x.

When an alkane burns in a limited supply of air, incomplete combustion occurs.

Explain why incomplete combustion of an alkane could be harmful to humans.

Titanium is an important metal in industry.

Titanium dioxide (TiO₂) can be converted into titanium metal in two stages.

Stage 1 titanium dioxide is converted into titanium(IV) chloride (TiCl₄)

Stage 2 titanium(IV) chloride is converted into titanium

This is the equation for the reaction in stage 1.

Calculate the volume, in dm³, of chlorine gas at rtp needed to react completely with 20 tonnes of titanium dioxide.

Give your answer in standard form.

[1 tonne = 10⁶g M_r of TiO₂ = 80]

[molar volume of chlorine gas at rtp = 24 dm³]

In stage 2, titanium(IV) chloride vapour is passed through molten magnesium in a container filled with argon.

This is the equation for the reaction in stage 2.

Explain why the container is filled with argon rather than air.

Aeroplanes are made of an alloy containing aluminium and titanium.

Explain why the alloy is stronger than pure titanium metal.

You may include diagrams in your answer.

Complete the table to show the relative mass and relative charge of a proton and a neutron.

Magnesium has three isotopes.

State the meaning of the term isotopes.

The symbol for an atom of one isotope of magnesium is

Give the number of protons, neutrons and electrons in one atom of this isotope.

number of protons

number of neutrons

number of electrons

A sample of magnesium contains these percentages of the three isotopes.

Mg‑24 = 79.00 % Mg‑25 = 10.00 % Mg‑26 = 11.00 %

Use this information to show that the relative atomic mass of magnesium is 24.32

One mole of magnesium has a mass of 24.32 g. There are 6.022 × 10²³ atoms in one mole.

Calculate the mass, in grams, of one atom of magnesium.

Give your answer to 4 significant figures.

The equation for the reaction between magnesium and oxygen is

Determine the maximum amount, in moles, of magnesium oxide that can be produced from 0.50 mol of magnesium and 0.20 mol of oxygen.

A student uses this apparatus to investigate the rate of reaction between calcium carbonate chips and dilute hydrochloric acid.

Every 20 seconds the student records the reading on the balance.

Explain why using a cotton wool plug increases the accuracy of the student’s results.

Complete the equation for the reaction by adding the state symbols.

The student uses the balance readings to find the decrease in mass of the flask and contents.

The graph shows the student’s results.

Give a reason why there are some calcium carbonate chips remaining in the flask when the reaction stops.

State how the student would know when the reaction has stopped.

Use the graph to determine the amount, in moles, of carbon dioxide produced during the reaction.

[M_r of CO₂ = 44]

Use the graph to calculate the rate of reaction, in grams per second, at time 60 seconds.

Show your working on the graph.

The student repeats the investigation by diluting the original hydrochloric acid.

The student then determines the initial rate of reaction at different percentage concentrations of the original hydrochloric acid.

The graph shows the student’s results.

Describe the relationship between the initial rate of reaction and percentage concentration of the original hydrochloric acid.

Explain why changing the concentration of hydrochloric acid has an effect on the initial rate of reaction.

This question is about the reactions of Group 1 metals with water.

A teacher adds a piece of sodium to some water containing universal indicator.

The equation for this reaction is

The sodium floats on the surface of the water and the universal indicator changes colour because an alkaline solution is formed.

Give two other observations.

Give the final colour of the universal indicator.

The diagram represents an atom of lithium and an atom of sodium.

Give a reason why lithium and sodium have similar reactions with water.

Explain why lithium is less reactive than sodium.

The teacher adds 0.150 g of lithium to an excess of water and collects the hydrogen gas produced.

The equation for the reaction is

The teacher collects 254 cm³ of hydrogen gas at room temperature and pressure (rtp).

Show by calculation that 1 mol of hydrogen gas has a volume of approximately 24 000 cm³ at rtp.

This question is about three stages in the manufacture of sulfuric acid.

In stage 1, sulfur is burned in oxygen to form sulfur dioxide gas.

State one environmental problem caused by the release of sulfur dioxide into the atmosphere.

A mass of 6.4 tonnes of sulfur is burned to produce sulfur dioxide gas.

Calculate the maximum volume, in dm³, of sulfur dioxide gas that can be produced at rtp.

[molar volume of sulfur dioxide gas at rtp = 24 dm³]

[1 tonne = 10⁶g]

Give your answer in standard form.

In stage 2, sulfur dioxide is reacted with oxygen to form sulfur trioxide gas.

The yield of sulfur trioxide is approximately 98%.

A catalyst is used in this reaction.

Explain how a catalyst increases the rate of a reaction.

The temperature is kept constant.

Give a reason why increasing the pressure would increase the yield of sulfur trioxide.

Suggest why it is not necessary to increase the pressure in stage 2.

In stage 3, the sulfur trioxide is reacted with concentrated sulfuric acid to form a liquid called oleum, H₂S₂O₇

The oleum is then added to water to form concentrated sulfuric acid.

Complete the chemical equations for these two reactions.

Sulfuric acid reacts with ammonia to form ammonium sulfate, (NH₄)₂SO₄

Calculate the percentage by mass of nitrogen in ammonium sulfate.

[M_r of (NH₄)₂SO₄ = 132]

A sample of a gaseous hydrocarbon, X, has a volume of 600 cm³ at room temperature and pressure (rtp).

Calculate the amount, in moles, of hydrocarbon X in the sample.

[molar volume of a gas = 24 000 cm³ at rtp]

The mass of the sample of hydrocarbon X is 1.45 g.

Show that the relative molecular mass (M_r) of X is 58

Hydrocarbon X is an alkane.

Show that the molecular formula of X is C₄H₁₀

Give the displayed formula of the branched‑chain isomer of hydrocarbon X.

A student uses the reaction between zinc and dilute sulfuric acid to prepare some zinc sulfate crystals.

Complete the equation for this reaction by giving the correct state symbols.

State what would be observed during this reaction.

The student adds excess zinc to a beaker of dilute sulfuric acid.

Explain why it is necessary to add excess zinc.

Draw a diagram of the apparatus the student should use to remove the unreacted zinc and collect the zinc sulfate solution.

The student obtains a pure, dry sample of zinc sulfate crystals.

The formula of zinc sulfate crystals is ZnSO₄.7H₂O

Calculate the relative molecular mass (M_r) of zinc sulfate crystals.

The student uses 0.0200 mol of dilute sulfuric acid in her preparation.

Show that the maximum mass of zinc sulfate crystals that the student could obtain is about 6 g.

The student obtains a mass of 4.28 g of zinc sulfate crystals.

Calculate the percentage yield of the zinc sulfate crystals.

Give your answer to three significant figures.

This question is about the reduction of metal oxides.

Solid oxides of copper can be reduced by reacting them with methane gas.

Complete the equation for the reaction between copper(II) oxide and methane.

Include state symbols.

_........ CuO(_........ ) +_........ CH₄(_........ ) _........ Cu(_........ ) +_........ CO₂(_........ ) +_........ H₂O(_........

A teacher uses this apparatus to demonstrate the reaction between a different oxide of copper and methane.

The teacher heats the oxide of copper until the reaction is complete.

The table shows the teacher’s results.

Use the teacher’s results to show that the empirical formula of this oxide of copper is Cu₂O

The teacher wears safety glasses and a lab coat during the demonstration.

Give one other safety precaution that she should take.

Iron forms when iron(III) oxide is heated with carbon.

The equation for the reaction is

State how the equation shows that iron(III) oxide is reduced.

State why carbon monoxide should not be released into the atmosphere.

Calculate the maximum mass, in tonnes, of iron that can be produced when 30.0 tonnes of iron(III) oxide are reacted with an excess of carbon.

[1 tonne = 1.0 × 10⁶ g]

A mixture of 25 000 mol of iron(III) oxide and 840 000 g of carbon is heated.

Use this equation to show that the iron(III) oxide is in excess.

This question is about Group 1 metals and their reactions.

When lithium is added to water, bubbles of hydrogen gas are observed.

Give two other observations that could be made.

Give the test for hydrogen gas.

Give one observation that would be different if potassium is used instead of lithium.

The diagram represents an atom of lithium and an atom of potassium.

Explain why potassium is more reactive than lithium.

The equation for the reaction between lithium and water is

A mass of 0.500 g of lithium reacts with an excess of water.

Calculate the volume, in cm³, of hydrogen gas produced at rtp.

[molar volume of a gas at rtp = 24 000 cm³]

Give your answer to three significant figures.

In a reaction between lithium and water, 150 cm³ of lithium hydroxide solution is formed.

The lithium hydroxide solution is then completely neutralised by 24.85 cm³ of 0.100 mol/dm³ sulfuric acid.

The equation for the neutralisation is

Calculate the concentration, in mol/dm³, of the lithium hydroxide solution.

concentration =_........ mol/dm