Redox (Cambridge (CIE) IGCSE Chemistry)

Separate: Chemistry and Extended Only

The following reactivity series shows both familiar and unfamiliar elements in order of decreasing reactivity. Each element is represented by a redox equation.

Rb Rb⁺ + e^–       
Mg  Mg²⁺ + 2e^–       
Mn Mn²⁺ + 2e^–       
Zn  Zn²⁺ + 2e^–       
H₂  2H⁺ + 2e^–       
Cu  Cu²⁺ + 2e^–       
Hg  Hg²⁺ + 2e^–

Two of the uses of the series are to predict the thermal stability of compounds of the metals and to explain their redox reactions.

i) Define in terms of electron transfer the term oxidation.

[1]

ii) Explain why the positive ions in the above equations are oxidising agents.

[1]

i) Which metals in the series above do not react with dilute acids to form hydrogen?

[1]

ii) Describe an experiment which would confirm the prediction made in (b)(i).

[1]

Separate: Chemistry and Extended Only

i) Which metal in the series above can form a negative ion which gives a pink / purple solution in water?

[1]

ii) Describe what you would observe when zinc, a reducing agent, is added to this pink / purple solution.

[1]

The table below shows the elements in the second period of the Periodic Table and some of their oxidation states in their most common compounds.

i) What does it mean when the only oxidation state of an element is zero?

[1]

ii) Explain why some elements have positive oxidation states but others have negative ones.

[2]

iii) Select two elements in the table which exist as diatomic molecules of the type X₂.

[2]

Separate: Chemistry and Extended Only

Beryllium hydroxide, a white solid, is an amphoteric hydroxide. 

i) Name another metal which has an amphoteric hydroxide.

[1]

ii) Suggest what you would observe when an excess of aqueous sodium hydroxide is added gradually to aqueous beryllium sulfate.

[2]

i) Give the formulae of lithium fluoride and nitrogen fluoride.

  Lithium fluoride ................................................................................................... Nitrogen fluoride ..................................................................................................

[2]

ii) Predict two differences in their properties.

[2]

iii) Explain why these two fluorides have different properties.

[2]

When the oxide, Co₃O₄  is heated in hydrogen, cobalt metal is formed.

Co₃O₄   +  4H₂   →   3Co   +   4H₂O

Explain how this equation shows that Co₃O₄ is reduced.

Separate: Chemistry and Extended Only

When copper(II) oxide is heated at 800 °C it undergoes the reaction shown by the  equation.

 Identify the changes in oxidation numbers of copper and oxygen in this reaction.

 Explain in terms of changes in oxidation numbers why this is a redox reaction.  

change in oxidation number of copper: from …………… to …………… change in oxidation number of oxygen: from …………… to ……………

explanation ...

The reaction of iron with steam is shown.

3Fe + 4H₂O → Fe₃O₄ + 4H₂

How does this equation show that iron gets oxidised?

Ammonia is manufactured by the Haber process.

Ammonia, NH₃, is used to produce nitric acid, HNO₃. This happens in a three-stage process.

Stage 1 is a redox reaction.

4NH₃ + 5O₂ → 4NO + 6H₂O

Identify what is oxidised in stage 1.

Give a reason for your answer.

substance oxidised .....................................

reason    .....................................

Zinc and copper are elements next to each other in the Periodic Table.

Aqueous potassium iodide reacts with aqueous copper(II) sulfate to produce iodine.

Balance the chemical equation for this reaction.

Extended Only

Deduce the charge on the copper ion in CuI.

Extended Only

In terms of electron transfer, explain why copper is reduced in this reaction.

Separate: Chemistry and Extended Only

Identify the reducing agent.

Separate: Chemistry and Extended Only

Sulfur dioxide is a reducing agent.

Give the colour change that occurs when excess sulfur dioxide is bubbled into acidified aqueous potassium manganate(VII).

starting colour of the solution ..............................................
final colour of the solution       ..............................................

Separate: Chemistry and Extended Only

Displacement reactions occur between metals and metal ions. Displacement reactions can be used to determine the order of reactivity of metals such as lead (Pb), nickel (Ni), and silver (Ag). The ionic equation for a displacement reaction is shown.

The ionic half-equations for this reaction are shown.

The ionic half-equations show that electrons are donated by nickel atoms and accepted by lead ions. Identify the reducing agent in the displacement reaction. Give a reason for your answer. reducing agent......................................... reason              ..........................................

Extended Only

What is the general term given to the type of reaction in which electrons are transferred from one species to another?

Vanadium is a transition element. It has more than one oxidation state.
The element and its compounds are often used as catalysts.

Predict three physical properties of vanadium which are typical of transition elements.

Separate: Chemistry and Extended Only

Vanadium(V) oxide is used to catalyse the exothermic reaction between sulfur dioxide and oxygen in the Contact Process.

2SO₂ + O₂ 2SO₃

The rate of this reaction can be increased either by using a catalyst or by increasing the temperature. Explain why a catalyst is used and not a higher temperature.

Separate: Chemistry and Extended Only

The oxidation states of vanadium in its compounds are V(+5), V(+4), V(+3) and V(+2).

The vanadium(III) ion can behave as a reducing agent or an oxidising agent.

i) Indicate on the following equation which reactant is the oxidising agent.

2V³⁺ + Zn → 2V²⁺ + Zn²⁺

[1]

ii) Which change in the following equation is oxidation? Explain your choice.

V³⁺ + Fe³⁺ → V⁴⁺ + Fe²⁺

[2]

The following are examples of redox reactions.

Bromine water was added to aqueous sodium sulfide. The ionic equation is:

Br₂ (aq) + S^2– (aq) → 2Br^– (aq) + S(s)

i) Describe what you would observe when this reaction occurs.

[2]

ii) Write a full symbol equation for this reaction.

[1]

iii) Explain, in terms of electron transfer, why bromine is the oxidant (oxidising agent) in this reaction.

[2]

Extended Only

Iron and steel in the presence of water and oxygen form rust.

The reactions involved are:
reaction 1

Fe → Fe²⁺ + 2e^–

The electrons move through the iron onto the surface where a colourless gas forms.

reaction 2

Fe²⁺ + 2OH^– → Fe(OH)₂

reaction 3

..........Fe(OH)₂ + O₂ + ..........H₂O → ..........Fe(OH)₃

The water evaporates to leave rust.

i) What type of reaction is reaction 1?

[1]

ii) Deduce the name of the colourless gas mentioned in reaction 1.

[1]

iii) What is the name of the iron compound formed in reaction 2?

[1]

iv) Balance the equation for reaction 3. ..........Fe(OH)₂ + O₂ + ..........H₂O → ..........Fe(OH)₃

[1]

v) Explain why the change Fe(OH)₂ to Fe(OH)₃ is oxidation.

[1]

vi) Explain why iron in electrical contact with a piece of zinc does not rust.

[3]

The distinctive smell of the seaside was thought to be caused by ozone, O₃. Ozone is a form of the element oxygen.

A mixture of oxygen and ozone is formed by passing electric sparks through oxygen.

3O₂ 2O₃

Suggest a technique that might separate this mixture. Explain why this method separates the two forms of oxygen.

Extended Only

Ozone is an oxidant. It can oxidise an iodide ion to iodine.

2I^– + O₃ + 2H⁺ → I₂ + O₂ + H₂O

i) What would you see when ozone is bubbled through aqueous acidified potassium iodide?

[2]

ii) Explain in terms of electron transfer why the change from iodide ions to iodine molecules is oxidation.

[1]

iii) Explain, using your answer to b(ii), why ozone is the oxidising agent in this reaction.

[1]

Extended Only

It is now known that the smell of the seaside is due to the chemical dimethyl sulfide, (CH₃)₂S.

i) Draw a diagram that shows the arrangement of the outer electrons in one molecule of this covalent compound. Use x to represent an electron from a carbon atom. Use ● to represent an electron from a hydrogen atom. Use o to represent an electron from a sulfur atom.

[3]

ii) Name the three compounds formed when dimethyl sulfide is burnt in excess oxygen.

[2]

An ore of tungsten conatins WO₃. Tungsten can be obtained from WO₃ by reacting it with hydrogen. The reaction is a redox reaction, water is also produced.

Write the balanced symbol equation for the reaction.

Extended Only

Explain reduction and oxidation in terms of electrons and oxidation numbers.

Separate: Chemistry and Extended Only

The oxidation number of oxygen is -2 and this remains unchanged during the reaction.

Identify, using oxidation numbers, what has been oxidised and what has been reduced in this reaction.

Extended Only

50 tonnes of the tungsten ore were processed. The ore was found to contain 3.4% of WO₃ by mass. Calculate the maximum mass, in tonnes, of tungsten that could be obtained from this ore.

Give your answer to three significant figures.

[A_r (W) = 184;    M_r (WO₃) = 232]

maximum mass of tungsten = ................................................. tonnes

Separate: Chemistry and Extended Only

What is meant by the term oxidising agent?

Separate: Chemistry and Extended Only

The dichromate ion, Cr₂O₇^2-, reacts with sulfite ions, SO₃^2-, according to the following equation 

Cr₂O₇^2- + 8H⁺ + 3SO₃^2-  →  2Cr³⁺ + 4H₂O + 3SO₄^2- 

Are the sulfite ions, SO₃^2-, acting as an oxidising or reducing agent. Justify your answer.

Separate: Chemistry and Extended Only

Identify which species is acting as the oxidising agent in the following reaction. Justify your answer.

Cl₂ + 2Br^– →  2Cl^– + Br₂ 

Separate: Chemistry and Extended Only

Redox titrations can be carried out between reducing agents and oxidising agents in a similar way to the method used for titrations between acids and alkalis.

Suggest why an indicator is not needed when potassium manganate(VII) is added to a reducing agent.