Modul Kimia

Tingkatan 5

Chapter 12 Oxidation and Reduction

Concept Map

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Oxidation and Reduction

Oxidising agent-electron gain

Reducing agent- electron loss

Oxidation process- gain of oxygen- increase in oxidation number- loss of hydrogen- loss of electron

Reduction process- loss of oxygen- decrease in oxidation number- gain of hydrogen- gain of electron

Redox Reactions

Combustion of metal in oxygen or chlorine

Reactivity series

Prevention of corrosion

Displacement of metals from their salts

Electrochemical series

Electrolytic cells

Displacement of halogens from halides

Heating metal oxides with carbon or hydrogen

Oxidising Power of halogen decreases down Group 17

Position of C and H in reactivity series of metals

Extraction of metals

12.1 Redox Reactions

1. Oxidation is defined as

(a) the addition of oxygen to a substance

(b) the elimination of hydrogen from a substance

(c) an increase in the oxidation number of an atom

2. Reduction is defined as

(a) the elimination of oxygen from a substance

(b) the addition of hydrogen from a substance

(c) a decrease in the oxidation of an atom

Gimmick

The first chemist who explained the oxidation and reduction reactions was Antoine Lavoisier.

3. An oxidising agent is a substance which oxidises another substance and itself is being reduced.

4. A reducing agent is a substance which reduces another substance and itself is being oxidised.

5. In terms of electron transfer,

(a) oxidation is the process of electron loss

(b) reduction is the process of electron gain

6. During a redox reaction, an oxidising agent gains electron from the reducing agent, while a reducing agent donates electron to

the oxidising agent.

Gimmick

An example of redox reaction that occurs in our everyday life is photosynthesis. During photosynthesis, carbon dioxide combines

with water to form sugar and oxygen.

6CO2 + 6H2O → C6H12O6 + 6O2

Carbon dioxide is reduced while water is oxidized.

7. In general, metals are reducing agents as they tend to lose their valence electrons during chemical reactions. On the other

hand, non-metals are generally oxidising agents as they tend to gain electrons during chemical reactions.

8. In a redox, oxidation and reduction processes occur simultaneously.

12.2 Oxidation Number

1. Oxidation number is the charge carried by an ion or an atom of an element in a compound.

2. The oxidation number has the same value as the valency of an element or ion but with positive(+) or negative(-) sign affixed to

it.

3. Oxidation number of an element is zero.

Element Oxidation Number

Na 0

H2 0

Cl2 0

S8 0

P4 0

4. The oxidation number of hydrogen is always +1 except in hydrides of metals where it is –1.

Compound Oxidation Number of

H atom

H2O +1

CH4 +1

H2S +1

NaH -1

MgH2 -1

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5. The oxidation number of oxygen is always –2 except in peroxides where it is –1.

Compound Oxidation Number of

O atm

H2O -2

SO2 -2

HNO3 -2

H2O2 -1

BaO2 -1

6. For monoatomic ions, the oxidation number of the atom is equal to the ionic charge.

Ion Oxidation Number of

element

Ag + +1

Zn 2+ +2

Cr 3+ +3

Br – -1

O 2 – -2

Gimmick

Some times the oxidation number of an element in a compound can be zero. One good example is 1,1,1-trichloroethane, C2H3Cl3..

In this molecule, the oxidation numbers of H, Cl, and C are +1, -1, and 0 respectively.

7. For polyatomic ions, the sum of the oxidation numbers of each atom is equal to the net charge of the ion.

Ion Oxidation number of

underlined atom

MnO 4 – x + 4(-2) = -1

x = +7

S 2 O 3 2 – 2x + 3(-2) = -2

x = +2

NH 4 + x + 4(+1) = +1

X = -3

SO 4 2– X + 4(-2) = -2

X = +6

8. Naming compounds using IUPAC system:

Formula Oxidation

number of metal

Name of compound

CuCl +1 Copper(I) chloride

CuBr2 +2 Copper(II) bromide

FeCl2 +2 Iron(II) chloride

FeCl3 +3 Iron(III) chloride

MnO2 +4 Manganese(IV) oxide

KMnO 4 +7 Potassium manganate(VII)

PbCl2 +2 Lead(II) chloride

Fe2O3 +3 Iron(III) oxide

Gimmick

Oil paintings are darkened over the years due to the reaction between lead compounds and hydrogen sulphide gas in the air.

Pb 2+ + H2S PbS + 2H +

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Careful treatment of old paintings with hydrogen peroxide solution can restore its original colour. This is because lead(II)

sulphide is oxidized to lead(II) sulphate.

PbS + 4H2O2 PbSO4 + 4H2O

12.3 Analysis of Redox Reactions

1. Addition of oxygen

2CuO + C 2Cu + CO2

CuO loses an oxygen atom and hence it is reduced.

C gains an oxygen atom and hence it is oxidised.

2. Addition of hydrogen

H2S + Cl2 2HCl + S

H2S loses hydrogen atom and hence it is oxidised.

Cl2 gains hydrogen atoms and hence it is reduced.

3. Change in oxidation number

2FeCl2 + Cl2 2FeCl3

+2 +3

0 -1

Oxidation number of iron increases from +2 to +3, hence FeCl2 is oxidised.

4. Transfer of electrons

2Na + Cl2 2NaCl

Na loses one valence electron to form Na +, hence sodium is oxidised.

Cl2 gains two electrons to form Cl –, hence chlorine is reduced.

12.4 Displacement Reactions

1. Moving down Group 17, strength of halogen as an oxidising agent decreases.

F2 > Cl2 > Br2 > I2

2. The tendency of halogens to be reduced decreases when moving down the group.

3. A halogen occupying higher position in the group can oxidise a halide which occupies a lower position in the group.

4. Thus chlorine oxidises bromine and iodide to iodine. But bromine cannot oxidise chloride to chlorine.

Cl2 + 2KBr 2KCl + Br2

Br2 + 2KCl 2KBr + I 2

5. Metals are arranged in Electrochemical Series as follows:

K K+ + e

Na Na + + e

Ca Ca 2+ + 2e Tendency to form

Tendency to Mg Mg 2+ + 2e metals,

donate electrons, Al Al 3+ + 3e oxidising power increases

reducing power Zn Zn 2+ + 2e

increases Fe Fe 2+ + 2e

Sn Sn 2+ + 2e

Pb Pb 2+ + 2e

H H + + e

Cu Cu 2+ + 2e

Ag Ag+ + e

6. When zinc is added to an aqueous solution of copper(II) sulphate, the blue solution is decolourized. A reddish-brown solid is

formed. Zinc is more electropositive than copper. Zinc displaces copper from its aqueous solution.

Zn + Cu 2+ Zn 2+ + Cu

7. Zinc reduces Cu 2+ to copper metal while itself is oxidised by Cu 2+ to Zn 2+ .

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12.5 Conversion of iron(II) to iron(III) and Vice Versa

1. Aqueous solution of Fe 2+ can be oxidized to Fe 3+ by treating it with the following oxidizing agents:

(a) chlorine (b) bromine (c) acidified potassium manganate(VII) (d) acidified potassium dichromate(VI)

2. Aqueous solution of iron(II) is light green while iron(III) is yellow.

3. Equations:

Cl2 + 2 Fe 2+ 2 Fe 3+ + 2Cl –

Br2 + 2 Fe 2+ 2 Fe 3+ + 2Br –

MnO 4 – + 5Fe 2+ + 8 H + Mn 2+ + 5Fe 3+ + 4H2O

Cr 2 O 7 2 – + 6Fe 2+ + 14H + 2Cr 3+ + 6Fe 3+ + 7H2O

4. Aqueous solution of Fe 3+ can be oxidized to Fe 2+ by treating it with the following reducing agents:

(a) aqueous iodide from KI

(b) zinc metal

(c) sulphur dioxide

5. Equations:

2I – + 2 Fe 3+ 2 Fe 2+ + I2

Zn + 2 Fe 3+ 2 Fe 2+ + Zn 2+

SO2 + 2H2O + 2 Fe 3+ 2 Fe 2+ + 2 H + + H2SO4

6. In the reaction between aqueous KI and iron(III), a brown precipitate of iodine is produced.

Quick Test

1. Determine the oxidation numbers of the underlined elements.

(a) K2S2O3 (b) Al2(SO4)3 (c) NaClO3 (d) KMnO4

(e) H2SO3 (f) NO3 - (g) ClO2 -

2. Identify the oxidizing and reducing agents in the following table.

Reaction Reducing agent Oxidising agent

H2O2 + 2I - + 2H +

I2 + 2H2O

Cu + 2AgNO3 Cu(NO3)2 + 2Ag

Br2 + 2 Fe 2+ 2 Fe 3+

+ 2Br -

Zn + 2HCl ZnCl2 + H2

CuO + H2SO4 CuSO4

+ H2O

3. Based on the following equation

2MnO4 - + 16H + + 10I - 2Mn 2+ + 5I2 + 8H2O

explain the above redox reaction in terms of electron transfer and change in oxidation numbers.

4. Mg + Cu 2+ Cu + Mg 2+

(a) What would you observe when a piece of magnesium is added to copper(II) sulphate solution?

(b) Write half-equations to represent the changes.

5. Write the formula of the following compounds.

(a) iron(III) sulphate (b) magnesium chlorate(V)

(c) copper(I) oxide (d) potassium dichromate(VI)

(e) sodium chromate(VI) (f) potassium hexacyanoferrate(III)

(g) cobalt(II) oxide

12.6 Transfer of Electrons at a Distance

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1. A redox reaction occurs when solutions of an oxidizing agent and a reducing agent are separated by an electrolyte in a U-tube

and connected by electrical wires.

2. When the circuit is completed, electrons are transferred from the negative electrode(anode) to the positive electrode(cathode)

through the external wire.

3. The electrode that is immersed in the reducing agent is the negative terminal where oxidation occurs.

4. The electrode that is immersed in the oxidizing agent is the positive terminal where reduction occurs.

5. The electrolyte, usually dilute sulphuric acid, serves as a salt bridge, which functions as follows:

(a) to separate the oxidizing agent from the reducing agent

(b) to complete the circuit so that ions can flow through it.

6. Other electrolytes which can be used as a salt bridge include

(a) potassium nitrate solution (b) potassium chloride solution.

7. Redox reaction between bromine and iron(II) sulphate:

Reaction at the negative terminal:

Fe 2+ ions are oxidized to Fe 3+. The colour of solution changes from light green to yellow.

Fe 2+ Fe 3+ + e

The electrons released by Fe 2+ travels through the external wire to the cathode.

Reaction at the cathode:

Bromine molecules gain electrons and being reduced to bromide. The reddish-brown colour of bromine solution disappears.

Bromine acts as the oxidizing agent.

Br2 + 2e 2 Br –

Overall equation: Br2 + 2 Fe 2+ 2 Fe 3+ + 2Br –

12.7 Rusting As a Redox Reaction

1. The corrosion or rusting of metals involves redox reactions. It is an electrochemical process.

2. Corrosion is the reaction between a metal and oxygen to form layer of metallic oxide.

3. Corrosion requires oxygen and water.

(a) An electrochemical cell is formed when iron comes into contact with water.

(b) Oxidation occurs at the anode region of the metal surface. Here iron atoms lose electrons to form ions Fe 2+.

Fe Fe 2+ + 2e

(c) The electrons then flow through the metal to the cathode region where hydroxide ions are formed by reduction.

O2 + 2H2O + 4e 4 OH –

The hydroxide ions dissolves in the water.

(d) The Fe 2+ ions then combine with the hydroxide ions to form a green precipitate of iron(II) hydroxide.

Fe 2+ + 2 OH – Fe(OH)2

(e) The iron(II) hydroxide is oxidized further by oxygen to form a brown solid of iron(III) hydroxide. Finally the iron(III)

hydroxide decomposes to form a brown solid called rust,

Fe2O3.xH2O.

5. The presence of dilute acids or salt solutions increases the rate of corrosion of metals.

6. Iron rusts faster in industrials areas because of acidic gases such as sulphur dioxide.

12.8 Reactivity of Metals towards Oxygen

1. Metals are oxidized by oxygen when they are heated in the presence of air. The final product is the oxide of the metal.

2. Metals show different reactivity towards oxygen in their reactions. More reactive metals react vigorously with oxygen while

the less reactive ones reacts slowly.

Gimmick

Gold is a very inert metal and is placed below silver in the activity series towards oxygen. The golden mask of King

Tutankhamun was discovered, the gold is still pure and uncorroded despite the fact that is has been buried for more than 3000

years.

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3. Oxygen gas can be prepared in the laboratory by heating solid potassium manganate(VII), potassium nitrate, and potassium

chlorate in the presence of manganese(IV) oxide.

2KMnO4 K2MnO4 + MnO2 + O2

2KNO3 KNO2 + O2

2KClO3 2KCl + 3O2

4. Reactivity series of metals:

Most reactive K Forms ions more easily

Na

Ca

Mg

Reactivity Al

decreases Zn

Fe

Sn

Pb

H

Cu

Least Ag

reactive

5. A metal can displace a less reactive metal from its oxide when the two are heated together.

Mg + ZnO MgO + Zn

Quick Test

1. Explain what happens when the following mixtures are heated.

(a) C + ZnO (b) Mg + CO2

(c) C + Al2O3

2. Complete the following table.

Electrolytic cell Chemical cell

Structure of cell

Energy change

Transfer of electrons

Type of reaction at anode

Type of reaction at cathode

6. The above displacement reaction is indeed a redox reaction. Mg reduces ZnO to Zn while itself is oxidized to MgO.

7. Hydrogen occupies the position between lead and copper. Hence hydrogen can reduce CuO to copper metal when the gas is

passed over hot copper(II) oxide.

H2 + CuO H2O + Cu

Assessment 12

Objective Questions

1. Which of the following statements is correct?

A. A reducing agent is a substance that receives electrons. B. Chlorine is a strong reducing agent.

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C. Reduction occurs at the cathode of an electrolytic cell. D. An oxidizing agent is a substance that releases electrons.

2. When potassium bromide solution is added to aqueous chlorine,

A. the solution changes from yellow to purple. B. bromine is displaced by chloride ions.

C. bromide ion is oxidized to bromine by chlorine. D. electrons are transferred from chlorine to bromide ions.

3. In which compound(s) would you find iron occurring with an oxidation number of +3?

I K3Fe(CN)6 II K2Fe(CN)6 III Fe2(SO4)3 IV Fe(NO3)2

A. III only B. I and III only C. II and IV only D. I, II, and III only

4. In a redox cell, a salt bridge between the two reagents serves to

I separate the oxidizing agent from the reducing agent

II allow electrons to flow from the negative terminal III cause a deflection in the galnometer

IV allow ions to flow between the two solutions in the U-tube

A. I and III only B. II and IV only C. I, II and III only D. I, II, and IV only

5. Which of the following conversions would involve a gain in electrons?

I Hydrogen atoms are converted into hydrogen ions. II Iodine atoms are converted into iodide ions.

III Manganate(VII) ions are converted into manganese(II) ions

IV Iron(III) ions are converted into iron(II) ions.

A. I and IV only B. II and III only C. I, II, and III only D. II, III, and IV only.

6. Aluminium can reduce the oxide of metal X but cannot reduce the oxides of metal Y and metal Z. Metal Z cannot reduce the

oxide of metal Y. Of the following, which one represents the arrangement of the four metals in descending order of their

reactivity?

A. Y, Al, X, Z B. Y, Z, Al, X C. X, Al, Z, Y D. Al, X, Z, Y

7. A mixture of aluminium oxide, copper(II) oxide and excess zinc is heated strongly until there is no further change. Which of

the following substances would be present in the final mixture?

I Aluminium II zinc oxide III Zinc IV Copper

A. I and IV only B. II and IV only C. I, II, and IV only D. I, II, III, and IV

8. Aluminium cannot be extracted from bauxite by heating it with coke because

A. aluminium is more reactive than oxygen B. the reaction between bauxite and coke is too vigorous

C. aluminium is more reactive than carbon D. aluminium is less reactive than carbon

Structured Questions

1. The figure below shows an experiment carried out to determine the position of hydrogen in the reactivity series using two

metallic oxides.

The oxide of P is heated strongly. Then the experiment is repeated by using the oxide of Q. The following results were obtained.

Substance Observation

Oxide of P Glows in red

Changes colour from reddish-brown to grey

Grey product is attracted by a magnet

Oxide of Q White solid does not glow

Turns yellow when hot; changes back to white

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when cold

(a) State two safety measures which you would take while doing the experiment. [2 marks]

(b) How would you prepare dry hydrogen in the laboratory? Draw a diagram to show it. [2 marks]

(c) From the results given in the above table, arrange hydrogen and metals P and Q in descending order of their reactivity.

[1 mark]

(d) Identify the metals P and Q based on the results in the above table. Explain your answer. [2 marks]

(e) Write the chemical equation for the reaction between the oxide of P and hydrogen. [1 mark]

(f) Name the liquid R that is formed as shown in the above figure. State a confirmatory test [2 marks]

(g) Predict what would happen if copper(II) oxide is heated in the presence of hydrogen. Explain your answer.

[2 marks]

2. The experiment in the figure below is carried out to study redox reaction through the transfer of electrons at a distance.

(a) (i) Name the oxidising and reducing agents in the experiment. [2 marks]

(ii) Show the flow of electrons and the two terminals on the above figure. [2 marks]

(b) What is the function of dilute sulphuric acid? [2 marks]

(c) (i) What happens at electrode A after an hour? [1 mark]

(ii) Write the half-equation to show the reaction at A. What redox reaction occurs

here? [1 mark]

(d) After the experiment, a small amount of the solution is taken out from A and tested with

sodium hydroxide solution. What would you observe? [1 mark]

(e) Name another reagent that can be used to replace the iron(III) chloride. [1 mark]

(f) (i) What would you observe at B after an hour? [1 mark]

(ii) Name the product formed at B and give the half-equation to show the reaction. [1 mark]

(iii) Give a test to confirm the formation of the product stated in (f)(ii) above. [1 mark]

Essay Question

1. (a)

Explain the statement above in the following reactions. Write balanced equations for reactions.

(i) The reaction between iron(II) chloride and chlorine. Explain in terms of change in oxidation number.

(ii) The reaction between magnesium and copper(II) oxide. Explain in terms of electron transfer.

(iii) The reaction between hydrogen sulphide and chlorine. Explain in terms of addition and loss of hydrogen.

[12 marks]

(b) Explain the rusting of iron in terms of redox reactions. [8 marks]

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A reduction is always accompanied by an oxidation reaction.

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