Reactivity of metals

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.

A student uses this apparatus to investigate the rate of reaction between dilute sulfuric acid and an excess of small pieces of zinc.

This is the student’s method.

Step 1 use 50 cm³ of dilute sulfuric acid

Step 2 add approximately 5 g of small zinc pieces to the sulfuric acid

Step 3 quickly connect the gas syringe

Step 4 record the reading on the gas syringe every 30 seconds until the reaction stops

Name a suitable piece of apparatus to measure the volume of sulfuric acid.

Give a reason why the mass of zinc pieces does not need to be measured accurately.

Give a reason why the student quickly connects the gas syringe in step 3.

State how the student would know when the reaction stops.

The graph shows the volume of gas collected in the syringe during the experiment.

A tangent to the curve has been drawn at a time of 80 s.

Use the tangent to calculate the rate of reaction at 80 s.

Show your working on the graph.

Give the unit.

Explain the shape of the graph in these regions.

from 0 s to 60 s

from 60 s to 150 s

from 150 s to 240 s

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.

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.

Iron metal was discovered several thousand years ago. Iron is produced industrially from iron(III) oxide by extraction using carbon.

Sodium is a metal in Group 1 of the Periodic Table. Sodium was discovered by Sir Humphrey Davy in 1807. It is produced industrially by the electrolysis of sodium chloride.

Explain, using the reactivity series, why carbon can be used to extract iron from iron(III) oxide.

Suggest why sodium was not discovered until 1807.

The diagram shows a simplified version of the Downs cell, which can be used to produce sodium industrially. The electrolyte is molten sodium chloride.

Explain why the molten electrolyte conducts electricity.

What is the formula of the cation in the electrolyte?

During the production of sodium, small explosions can be heard.

Explain why using sodium chloride that is not completely dry could cause these small explosions.

Give ionic half-equations for the reactions at the anode and cathode.

Explain why the reactivity of metals in Group 1 increases from lithium to potassium.

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.

Two substances are needed to cause iron to rust.

Name these two substances.

The box gives the names of some substances.

Use words from the box to answer these questions.

Give the name of a non-metallic element.

Give the name of a compound.

Give the name of the metal that is lowest in the reactivity series.

A student uses this apparatus to investigate the effect of heat on different solid metal carbonates.

This is the student’s method.

use a spatula to put some metal carbonate in the boiling tube

fit the delivery tube into position

pour some limewater into the test tube

start a timer and immediately begin to heat the metal carbonate

record the time when a change first occurs in the limewater

The student repeats the method using different metal carbonates.

When a metal carbonate is heated a reaction sometimes occurs.

The equation for the reaction is

metal carbonate metal oxide + carbon dioxide

State the name given to this type of reaction.

State two variables that the student should control in this investigation.

Suggest why bubbles appear in the limewater immediately after heating has started but before there is any change to the metal carbonate.

Explain the purpose of limewater in this investigation.

The table shows some of the results for the student’s investigation.

State the colour change that occurs for copper(II) carbonate.

Give a chemical equation for this reaction of copper(II) carbonate.

There is a relationship between the position of a metal in the reactivity series and how easily the metal carbonate reacts when heated.

Use the student’s results and your own knowledge to deduce this relationship.

State how you should extend the investigation to see if your deduction is correct.

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 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.

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 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.

A student uses this method to investigate the reaction of dilute hydrochloric acid with zinc.

● pour some dilute hydrochloric acid into a glass beaker

● record the initial temperature of the acid

● add a piece of zinc and stir the mixture

● record the temperature of the mixture after one minute

Write a word equation for the reaction of dilute hydrochloric acid with zinc.

The diagram shows the thermometer readings for this reaction.

Complete the table, giving all values to the nearest 0.1°C.

Another student repeats the method using five different metals to compare their reactivity.

This student uses a polystyrene cup instead of a glass beaker.

Explain why a polystyrene cup is better than a glass beaker in this investigation.

Give three factors that the student should keep constant in this investigation.

The table shows some of the student’s results.

State why there is no temperature change for copper.

Predict the temperature change for iron.

Deduce the order of reactivity of the five metals.

This question is about some Group 2 elements and their compounds.

Calcium reacts with water to produce calcium hydroxide and hydrogen gas.

Give the word equation for this reaction.

State two observations that would be made during this reaction.

Describe how a pure, dry sample of the insoluble salt, barium sulfate, could be made from the two solids sodium sulfate and barium chloride.

Give an ionic equation for the reaction that occurs.

Include state symbols in your equation.

When magnesium nitrate is heated, magnesium oxide, nitrogen dioxide and oxygen form.

The equation for the reaction is

What is the name for this type of reaction?

Calculate the total volume, in dm³, of gas produced at rtp when 7.7 g of magnesium nitrate completely reacts.

[Assume that the molar volume of a gas at rtp is 24 dm³]

[M_r of Mg(NO₃)₂ = 148]

Give your answer to two significant figures.