Oxidation & reduction in terms of electrons

When iron rusts, iron(III) oxide forms.

State two conditions needed for iron to rust.

What is the name of the process used to coat iron with zinc?

What is the correct order of reactivity of these four metals?

Describe a test to show that a solution contains Fe³⁺ ions.

test

result

A sample of iron(III) oxide reacts with excess carbon monoxide.

This is the equation for the reaction.

Explain why this is a redox reaction.

Calculate the minimum mass of iron(III) oxide needed to produce a theoretical yield of 28 g of iron.

[for Fe₂O₃ M_r = 160]

The actual yield of iron from this sample is 21 g.

Calculate the percentage yield of iron.

This question is about a compound called potassium permanganate.

A teacher makes a saturated solution of potassium permanganate using this method.

The teacher adds 25 cm³ of water at 50 °C to a beaker.

The teacher then adds solid potassium permanganate one spatula at a time, while stirring, until the solution is saturated.

Give the meaning of the term saturated solution.

Give a reason why crystals form when the saturated solution cools from 50 °C to room temperature.

The teacher pours 25 cm³ of a saturated solution of potassium permanganate at room temperature into an evaporating dish.

The teacher leaves the solution until all the water has evaporated.

The teacher records the mass of the evaporating dish and the mass of the evaporating dish with the dry crystals.

mass of evaporating dish = 64.10 g

mass of evaporating dish with dry crystals = 65.35 g

Calculate the solubility of potassium permanganate in grams per 100 g of water.

[1.00 cm³ of water has a mass of 1.00 g]

solubility =_........ g per 100 g of water

The equation shows the thermal decomposition of potassium permanganate.

2KMnO₄ K₂MnO₄ + MnO₂ + O₂

Describe a test to show that oxygen gas is produced.

Calculate the M_r of K₂MnO₄

Potassium permanganate is an oxidising agent.

Give a reason why potassium permanganate is an oxidising agent.

The table gives some information about five gases.

Use gases from the table to answer these questions.

You may use each gas once, more than once or not at all.

Name the gas needed for combustion.

Name the gas that is most abundant in air.

Name the gas that can bleach damp litmus paper.

Explain why hydrogen chloride is classified as a compound.

Why does chlorine have the highest boiling point of the gases in the table?

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]

This question is about magnesium metal and its compounds.

The diagram represents the structure of magnesium.

Complete the diagram by adding the labels.

Explain why magnesium is malleable.

Magnesium burns in oxygen gas to form solid magnesium oxide.

Complete the chemical equation for the reaction by balancing the equation and adding the state symbols.

Give a reason why magnesium is oxidised in this reaction.

A student uses this method to make a solution of magnesium chloride.

Step 1 measure 25 cm³ of dilute hydrochloric acid into a beaker

Step 2 add magnesium powder a little at a time

Step 3 keep adding magnesium powder until it is in excess

Step 4 remove the excess magnesium powder by filtration

This is the equation for the reaction.

Give a reason why the student uses magnesium powder rather than magnesium ribbon.

Give a reason why the student adds an excess of magnesium in step 3.

Calculate the minimum mass of magnesium needed to react with 25.0 cm³ of 2.00 mol / dm³ hydrochloric acid.

[for Mg, A_r = 24]

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.

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.

This question is about gases in the atmosphere.

A teacher uses this apparatus to determine the percentage of oxygen in air.

The teacher removes the stopper, ignites the magnesium ribbon and immediately replaces the stopper.

The magnesium reacts with oxygen to form magnesium oxide.

During the reaction the water level in the bell jar rises.

When the flame goes out, some magnesium remains in the basin.

Give the appearance of the magnesium oxide.

Give a chemical equation for the reaction of magnesium with oxygen.

Explain why the water in the bell jar rises.

The volume of air in the bell jar at the start of the reaction is 2000 cm³.

When the reaction ends, the apparatus cools down to room temperature.

Calculate the expected volume of gas in the bell jar at room temperature.

State why the gas remaining in the bell jar at the end of the reaction is approximately 99% nitrogen.

Sodium chloride is an ionic compound.

Explain why sodium chloride conducts electricity when it is molten or in solution, but not when it is solid.

A solution of sodium chloride can be electrolysed using this apparatus.

If the solution is dilute, a significant amount of oxygen collects at P.

Complete the ionic half-equation for this reaction.

If the solution is concentrated, chlorine is the main product that collects at P.

This is the ionic half-equation for the reaction.

State why this is an oxidation reaction.

The gas that collects at Q is hydrogen.

Give a test for hydrogen.

Explain how hydrogen gas forms at the negative electrode.

Hydrogen reacts with oxygen to form water.

The equation shows the covalent bonds in the molecules.

The table gives the bond energies.

Use the equation and the values in the table to calculate the enthalpy change, Δ H, for the reaction.

Include a sign in your answer.

Complete the diagram to show the energy levels of the reactants and products, and the enthalpy change, Δ H.

This question is about oxygen.

Oxygen can be stored in tanks as a liquid or as a gas.

Complete the diagram to show the arrangement of six more particles in a gas.

Give a reason why a tank can store much more oxygen as a liquid.

Identify a hazard when storing oxygen as a gas.

Sulfur burns in oxygen to form sulfur dioxide.

Give one observation that can be made when sulfur burns in oxygen.

Some universal indicator is added to sulfur dioxide.

Explain the final colour of the universal indicator.

This question is about the reactions of iron.

Iron rusts when exposed to water and oxygen.

Give the chemical name of the compound that forms when iron rusts.

What type of reaction occurs when iron rusts?

Galvanising is a method used to prevent iron from rusting.

Give the name of the metal used to galvanise iron.

When iron reacts with dilute sulfuric acid, the products are iron(II) sulfate and hydrogen.

Give a chemical equation for the reaction between iron and sulfuric acid.

Give a test for hydrogen.

An excess of iron is added to copper(II) sulfate solution.

Name the type of reaction that occurs.

State the appearance of the solid that forms in the reaction.

Give the reason why no reaction occurs when iron is added to magnesium sulfate solution.

This question is about iron.

One problem with iron is that it rusts.

Name the two substances that iron reacts with when it rusts.

State what type of reaction occurs when iron rusts.

Iron can be prevented from rusting by painting or by coating with zinc.

Explain how painting prevents iron from rusting.

Name the process used to coat iron with zinc.

Explain why a layer of zinc protects iron from rusting, even if the layer of zinc is scratched.

Iron is formed when aluminium reacts with iron(III) oxide.

This is the equation for the reaction.

Explain what this reaction shows about the relative reactivities of aluminium and iron.

Explain which substance acts as an oxidising agent in this reaction.

The box gives the names of some metals.

Use words from the box to answer these questions.

Identify the metal that burns with a bright white flame.

Explain which metal is most likely to be found in the Earth as an uncombined element.

Steel is an alloy of iron and carbon.

The diagram shows how the particles are arranged in steel.

State what is meant by the term alloy.

Explain why an alloy is less malleable than a pure metal.

This question is about the insoluble salt, lead(II) bromide.

Lead(II) bromide can be made by a precipitation reaction.

This is the equation for the reaction.

Pb(NO₃)₂(aq) + 2KBr(aq) PbBr₂(s) + 2KNO₃(aq)

Describe how solutions of lead nitrate and potassium bromide can be used to make a pure, dry sample of lead(II) bromide.

A solution containing 0.150 mol of lead(II) nitrate is reacted with an excess of potassium bromide solution.

A mass of 49.6 g of pure, dry lead(II) bromide is produced.

Show, by calculation, that the percentage yield of lead(II) bromide is 90.1%.

[for PbBr₂, M_r = 367]

A student investigates the change in electrical conductivity as dilute lead(II) nitrate solution is added to dilute potassium bromide solution.

This is the student’s method.

Step 1add 50 cm³ of potassium bromide solution to a beaker

Step 2 measure the electrical conductivity of the solution

Step 3add 10 cm³ of lead(II) nitrate solution to the beaker

Step 4 stir the mixture

Step 5 measure the electrical conductivity of the mixture

Repeat steps 3, 4 and 5 until a total of 50 cm³ of lead(II) nitrate solution has been added.

The table shows the student’s results.

Plot the student’s results on the grid.

Draw a line of best fit, ignoring the anomalous result.

Total volume of lead(II) nitrate in cm³

Explain the shape of the graph.

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

Further 10 cm³ volumes of lead(II) nitrate are added to the beaker so the lead(II) nitrate is in excess.

Predict what will happen to the conductivity of the mixture when the lead(II) nitrate is in excess.

The diagram shows the electrolysis of molten lead(II) bromide, PbBr₂

This is the ionic half‑equation for the formation of bromine at electrode A.

Give a reason why this half‑equation shows oxidation.

When copper(II) sulfate solution is electrolysed, copper forms at the negative electrode.

A student uses this apparatus to investigate the electrolysis of copper(II) sulfate solution.

Describe how the student could test a sample of copper(II) sulfate solution to show that it contains copper(II) ions.

Describe how copper metal forms at the negative electrode.

State the appearance of the copper that forms on the negative electrode.

Oxygen forms at the positive electrode.

Give a test for oxygen.

Complete the half-equation for the formation of oxygen at the positive electrode.

2H₂O _........ +_........ +_........

State why the formation of oxygen at the positive electrode is an oxidation reaction.

A student investigates the reactivities of four metals, aluminium, magnesium, copper and metal X.

The student adds pieces of magnesium ribbon to aqueous solutions of the sulfates of each metal.

After a few minutes the student removes the pieces of magnesium ribbon and records the appearance of each piece of magnesium.

Table 1 shows the student’s results.

Name the substance that causes the brown coating on the magnesium.

State why there is no change with magnesium sulfate solution.

The student repeats the experiment with pieces of metal X instead of pieces of magnesium.

Table 2 shows the student’s results.

Use the information from both tables to deduce the order of reactivity of aluminium, magnesium, copper and metal X.

Give a possible identity for metal X.

This ionic equation represents the reaction between magnesium and aluminium nitrate.

Explain, in terms of electrons, which species acts as a reducing agent in this reaction.

A teacher ignites a piece of magnesium ribbon and places it in a gas jar of oxygen.

This is the equation for the reaction.

Give two observations that would be made during this reaction.

The teacher demonstrates the reaction between aluminium and iron(III) oxide.

The teacher ignites the magnesium fuse. A very exothermic reaction then occurs between aluminium powder and iron(III) oxide.

The products of the reaction are iron and aluminium oxide.

State what is meant by the term exothermic.

Give a chemical equation for the reaction between aluminium and iron(III) oxide.

Explain why this reaction is a redox reaction.

A student uses this apparatus to determine the formula of magnesium oxide.

This is the student’s method.

• find mass of empty crucible and lid

• place some magnesium ribbon in crucible

• find mass of crucible, lid and magnesium ribbon

• heat crucible, lifting lid occasionally

• after five minutes, find mass of crucible, lid and magnesium oxide

Explain why the student lifts the lid.

The mass of magnesium oxide formed is less than expected.

Describe what the student should do next to obtain a more accurate value for the mass of magnesium oxide.

This question is about exothermic reactions.

A student uses this apparatus to measure the temperature increase when an excess of zinc powder is added to copper(II) sulfate solution.

Complete the word equation for the reaction.

zinc + copper(II) sulfate _........ +_........

Give a reason why the student uses a polystyrene cup inside a glass beaker.

State why zinc reacts with copper(II) sulfate solution.

The temperature at the start of the reaction is 19.7C.

The temperature at the end of the reaction is 48.3C.

Calculate the heat energy change, in joules, for the reaction.

[for the mixture, c = 4.2 J/g/C]

The reaction between zinc and silver nitrate solution is exothermic.

A mass of 0.65 g of zinc is added to excess silver nitrate solution.

The heat energy change is 800 J.

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

Include a sign in your answer.

ΔH =_........ kJ/mol

This is the ionic equation for the reaction between zinc and silver nitrate solution.

Explain, in terms of electrons, why this is a redox reaction.

This question is about the oxides of lead.

Yellow lead oxide (PbO) can be reacted with hydrogen to produce lead.

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

What is the charge on the lead ion in PbO?

Explain why the reaction of yellow lead oxide with hydrogen is a redox reaction.

Describe a physical test to show that the water produced in this reaction is pure.

When red lead oxide (Pb₃O₄) is heated, yellow lead oxide forms.

The equation for the reaction is

A scientist heats a known mass of red lead oxide in a crucible in a fume cupboard.

The scientist leaves the crucible to cool, then records the total mass of the crucible and its contents.

Describe what the scientist should do next to make sure that all the red lead oxide has reacted.

The red lead oxide used in the reaction has a mass of 5.48 g.

Calculate the maximum mass of yellow lead oxide that could form.

[M_r of Pb₃O₄ = 685 M_r of PbO = 223]

A potassium permanganate crystal is placed in a beaker of water.

After several days a coloured solution forms.

Give the names of the two processes that cause the coloured solution to form.

The formula of potassium permanganate is KMnO₄

How many different types of atom are in KMnO₄?

Calculate the relative formula mass (M_r) of KMnO₄

Potassium permanganate can be used as an oxidising agent.

State what is meant by the term oxidising agent.

This question is about magnesium and magnesium compounds.

Magnesium burns in oxygen to form magnesium oxide.

State two observations that would be seen during the reaction.

The diagram shows the electron configurations of a magnesium atom and an oxygen atom.

Describe the changes in the electronic configurations when magnesium reacts with oxygen to form the ionic compound magnesium oxide, MgO

Magnesium can be produced by the electrolysis of molten magnesium chloride.

State why magnesium cannot be produced by heating magnesium oxide with carbon.

Explain the different ways that magnesium and magnesium chloride conduct electricity.

magnesium

magnesium chloride

During the electrolysis of molten magnesium chloride, magnesium is formed at the negative electrode.

The ionic half‑equation for the reaction is

State why this is a reduction reaction.

Write an ionic half‑equation for the formation of chlorine at the positive electrode.

A teacher prepares the insoluble salt lead(II) bromide (PbBr₂) by mixing solutions of lead(II) nitrate and sodium bromide.

Describe what the teacher should do next to obtain a pure, dry sample of lead(II) bromide.

The teacher then sets up a circuit in a fume cupboard using the pure, dry sample of lead(II) bromide.

Explain why the lamp does not light when the lead(II) bromide is solid.

The teacher heats the lead(II) bromide.

When the lead(II) bromide is molten, the lamp lights and bromine forms at the positive electrode.

State what observation would be made at the positive electrode.

Explain how bromide ions in the molten lead(II) bromide become bromine molecules at the positive electrode.

Write an ionic half-equation for the reaction that occurs at the negative electrode.

Include state symbols in your equation.

Sulfur dioxide (SO₂) and hydrogen sulfide (H₂S) are both gases.

The two gases react together to form solid sulfur and water.

Complete the chemical equation for the reaction.

State why the sulfur dioxide is reduced in the reaction.

The diagram shows apparatus used to compare the speed at which particles of the two gases diffuse.

cotton wool soaked in pale yellow cotton wool soaked in sulfur dioxide solution solid hydrogen sulfide solution

The two pieces of cotton wool and rubber bungs are put in position at the same time.

A pale yellow solid soon forms.

Explain how the diagram shows that hydrogen sulfide gas diffuses more quickly than sulfur dioxide gas.

Deduce a relationship between the relative formula mass (M_r) of a gas and the speed at which a gas diffuses.

Use the A_r values to help you.

[A_r values: H = 1 S = 32 O = 16]

This question is about elements in Group 7 and their compounds.

The table gives information about some of these elements.

Predict the colour of fluorine at room temperature.

How many of the elements in the table are liquids at room temperature (20C)?

The element astatine is below iodine in Group 7.

Predict the formula of a molecule of astatine.

Sea water contains bromide ions.

Bromine can be obtained by bubbling chlorine through a sample of sea water.

The ionic equation for the reaction is

Explain which species acts as an oxidising agent in this reaction.

The reaction occurs because chlorine is more reactive than bromine.

Bromine is below chlorine in Group 7.

Explain the decrease in reactivity from chlorine to bromine.

Elements in Group 7 react with elements in Group 1 to form ionic compounds.

Which pair of ions both have the electronic configuration 2.8.8?

This question is about the metal aluminium.

Aluminium is malleable and conducts electricity.

The diagram shows the arrangement of the ions in aluminium metal.

Explain why aluminium is malleable.

Explain why aluminium conducts electricity.

Aluminium cannot be extracted by heating a mixture of carbon and aluminium oxide.

Give a reason why heating a mixture of aluminium oxide and carbon does not produce aluminium.

Aluminium is extracted industrially by the electrolysis of molten aluminium oxide Al₂O₃ at a temperature of about 950 °C.

Aluminium metal forms at the negative electrode and oxygen gas forms at the positive electrode. The positive and negative electrodes are made of graphite.

The diagram shows the apparatus used.

Explain how aluminium metal forms at the negative electrode.

Write an ionic half-equation for the formation of oxygen gas at the positive electrode.

Suggest why carbon dioxide gas is also produced at the positive electrode.

Aluminium reacts with iron(III) oxide. The reaction is exothermic.

The equation for the reaction is

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

Draw an energy level diagram for the reaction between aluminium and iron(III) oxide.

This question is about metals.

Metals can be arranged in a reactivity series based on their reactions with water and their reactions with dilute hydrochloric acid.

The table shows how four metals, P, Q, R and S, react with water and with dilute hydrochloric acid.

Identify which of the metals P, Q, R or S could be gold.

Suggest why the reaction between metal R and dilute hydrochloric acid was not done.

Use the information in the table to place the metals in order of reactivity from most reactive to least reactive.

Zinc is used to coat iron gates to prevent the iron from rusting.

State the name of this method of preventing iron from rusting.

State another method of preventing iron from rusting.

A mixture of zinc powder and copper(II) oxide is heated.

The chemical equation for the reaction that takes place is

State how the reaction shows that zinc is more reactive than copper.

Explain which substance is the oxidising agent.

A piece of magnesium ribbon is ignited and placed in a gas jar of oxygen.

The equation for the reaction is

Give two observations that would be made in this reaction.

State why this is an oxidation reaction.

A second piece of magnesium ribbon is ignited and placed in a gas jar of carbon dioxide.

A very exothermic reaction occurs, forming magnesium oxide and carbon.

State what is meant by the term exothermic.

Give the chemical equation for this reaction.

A fire starts in a warehouse where magnesium is stored.

Suggest why it would not be suitable to use a carbon dioxide fire extinguisher to put out this fire.

A student uses this apparatus to find the mass of magnesium oxide that forms when a known mass of magnesium is heated.

This is his method.

• find the mass of the crucible and lid

• place some magnesium ribbon in the crucible

• find the mass of the crucible, lid and magnesium

• heat the crucible with the lid on for a few minutes

• find the mass of the crucible, lid and magnesium oxide

Using this method, the mass of magnesium oxide formed is less than expected.

Explain two changes that the student should make to his method to obtain a mass of magnesium oxide closer to the expected mass.

The box gives the names of some metals.

Identify the metal from the box that burns with a bright white flame.

In the Earth, metals are found either in ores or as uncombined elements.

Explain which metal from the box is most likely to be found as an uncombined element.

This is the order of reactivity of four metals.

The method used to obtain a metal from its oxide depends on the reactivity of the metal.

Two possible methods are

Method 1 heating the metal oxide with carbon

Method 2 electrolysis

Explain which method should be used to obtain lead from lead(II) oxide, PbO

Include an equation for the formation of lead in your answer.

The diagram shows the arrangement of the particles in a pure metal.

Metals are often made into alloys to make them harder.

Explain why alloys are harder than pure metals.

Draw a diagram to support your answer.

This question is about the metal aluminium.

Draw a labelled diagram to represent the structure and bonding in a metal.

Explain why a metal conducts electricity.

Aluminium is used to make cans for drinks.

Give two properties of aluminium that make it suitable for this use.

Aluminium is extracted from aluminium oxide (Al₂O₃) by electrolysis.

The electrolyte is aluminium oxide dissolved in molten cryolite.

State why aluminium cannot be extracted by heating aluminium oxide with carbon.

Aluminium is produced at the negative electrode.

The ionic half-equation for the reaction is

State why this is a reduction reaction.

Complete the ionic half-equation for the reaction at the positive electrode.

This question is about gases in the atmosphere.

The box gives the names of some gases in the atmosphere.

Choose gases from the box to answer these questions.

Each gas may be used once, more than once or not at all.

Identify a noble gas.

Identify a gas that makes up about 78% of the atmosphere.

Identify a greenhouse gas.

Identify a gas produced by the thermal decomposition of calcium carbonate.

Sulfur reacts with oxygen to produce sulfur dioxide gas.

Write a chemical equation for this reaction.

State an environmental problem caused when sulfur dioxide gas dissolves in water in the atmosphere.

The diagram shows two samples of iron, A and B.

Sample B is coated with a thin layer of zinc.

Name the process used to coat iron with zinc.

The two samples of iron are left outside for several weeks.

A brown solid containing hydrated iron(III) oxide forms on sample A.

Give the common name for the brown solid.

Give the names of the two substances that react with the iron to form the brown solid.

Iron can be formed by reacting aluminium powder with iron(III) oxide.

The diagram shows how this reaction can be demonstrated.

When the magnesium fuse is lit, a very exothermic reaction occurs.

State the meaning of the term exothermic.

The equation for the reaction between aluminium and iron(III) oxide is

Explain what this reaction shows about the relative reactivities of aluminium and iron.

Explain why the reaction between aluminium and iron(III) oxide is a redox reaction.

The reaction between magnesium and copper(II) sulfate solution is exothermic.

This apparatus is used to measure the temperature increase when excess magnesium is added to 100 cm³ of copper(II) sulfate solution.

State why a reaction occurs when magnesium is added to copper(II) sulfate solution.

Complete the word equation for this reaction.

magnesium + copper(II) sulfate _........ +_........

The temperature at the start of the reaction is 20.2 °C.

The maximum temperature recorded is 56.3 °C.

Calculate the heat energy change, in joules, for the reaction.

[mass of 1.00 cm³ of solution = 1.00 g]

[c for the solution = 4.2 J/g/ °C]

Explain why it is better to use a polystyrene cup rather than a glass beaker in this experiment.

The reaction between zinc and copper(II) sulfate solution is also exothermic.

A mass of 0.500 g of zinc is reacted with an excess of copper(II) sulfate solution.

The heat energy change is 1.67 kJ.

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

Include a sign in your answer.

Give your answer to three significant figures.

H =_........ kJ/mol

The ionic equation for the reaction between zinc and copper(II) sulfate is

Explain why this is a redox reaction.

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 copper and its compounds.

Copper is a metal used for electrical wiring.

Explain why copper is a good conductor of electricity.

This apparatus is used to investigate the electrolysis of copper(II) sulfate solution with graphite electrodes.

Copper forms at the negative electrode and oxygen forms at the positive electrode.

State what would be observed at each electrode.

negative electrode_........

positive electrode_........

The ionic half‑equation for the reaction at the negative electrode is

State why this is a reduction reaction.

Explain why the copper(II) sulfate solution becomes paler blue during the electrolysis.

When hydrated copper(II) sulfate crystals are heated, anhydrous copper(II) sulfate forms.

A mass of 12.5 g of hydrated copper(II) sulfate crystals is heated in a crucible until all the water of crystallisation is removed.

A mass of 8.0 g of anhydrous copper(II) sulfate forms.

Show by calculation that the formula of hydrated copper(II) sulfate is CuSO₄.5H₂O

[M_r of CuSO₄ = 159.5 M_r of H₂O = 18]

This question is about the metal, lead.

Explain why metals, such as lead, are malleable.

A teacher uses this apparatus in a fume cupboard to demonstrate the electrolysis of lead(II) bromide.

The lead(II) bromide is heated until it melts.

When the lead(II) bromide melts, the lamp lights.

One of the products of this electrolysis is lead.

State why solid lead(II) bromide does not conduct electricity.

Bromine is formed by the oxidation of bromide ions at the positive electrode.

Complete the ionic half-equation for the oxidation of bromide ions.

Explain why lead metal forms at the negative electrode.

The teacher stops heating the mixture and allows it to solidify.

Suggest why the lamp stays alight.

This question is about the halogens and their compounds.

The table gives the colour and physical state at room temperature of the halogens.

Complete the table by predicting the colour of astatine and the physical state of fluorine at room temperature.

Chlorine gas is bubbled into a colourless solution of potassium bromide.

Explain why the solution turns orange.

Potassium bromide is an ionic compound.

Draw diagrams to show the outer electrons in a potassium ion and in a bromide ion.

Include the charges on the ions.

A student sets up a circuit to test the electrical conductivity of water, solid sodium chloride and aqueous sodium chloride.

The table shows the student’s results.

Explain these results, with reference to the structure and bonding of the substances.

A concentrated aqueous solution of sodium chloride is electrolysed using graphite electrodes.

Chlorine is formed at the positive electrode (anode).

Give an ionic half‑equation for the formation of chlorine at the positive electrode.

State why this ionic half‑equation represents an oxidation reaction.

Which substance is formed at the negative electrode (cathode)?