3 Oxidation and Reduction

Chapter 3

WAJA F5 Chemistry 2009Chapter 3 : Oxidation And Reduction

CHAPTER 3: OXIDATION AND REDUCTION

A.REDOX REACTIONS

Learning Outcomes:

You should be able to,

state what oxidation is

state what reduction is

explain what redox reaction is

state what oxidising agent is

state what reducing agent is

calculate the oxidation number of an element in a compound.

relate the oxidation number of an element to the name of its compound using the IUPAC

nomenclature.

explain with examples oxidation and reduction processes in terms of electron transfer

explain with examples oxidising and reducing agents in redox reactions.

write oxidation and reduction half-equations and ionic equations.

Activity 1

1)The meaning of oxidation and reduction in terms of:

(a)

Gain or Loss of Oxygen

Oxidation

Reduction

The process of gaining oxygen

The process of losing oxygen

Example :

Mg + PbO ( MgO + Pb

Mg gains oxygen to form MgO : The process is

PbO loses oxygen to form Pb : The process is ..

Mg is a reducing agent because

PbO is an oxidizing agent because ...

(b)

The loss and gain of hydrogen

Oxidation

Reduction

The process of losing hydrogen

The process of gaining hydrogen

Example :

H2S + Cl2 ( 2HCl + S

H2S loses hydrogen to form S : The process is ...

Cl2 gains hydrogen to form HCl : The process is .

H2S is a reducing agent because ..

Cl2 is an oxidizing agent because ..

(c)

the transfer of electrons

Oxidation

Reduction

The process of losing electrons

The process of gaining oxygen

Example :

2Na + Cl2 ( 2NaCl

Na ( Na+ + e // Cl2 + 2e ( 2Cl-

Na loses electron to form Na+ : The process is

Cl2 gains electrons to form 2Cl- : The process is ..

Na is a reducing agent because .

Cl2 is an oxidizing agent because ..

(d)

The change in oxidation number

Oxidation

Reduction

The increase in oxidation number

The decrease in oxidation number

Example :

2Na + Cl2 ( 2NaCl

Na ( Na+ // Cl2 ( 2Cl-

Oxidation number Na : 0 ( +1 // Cl : 0 ( -1

Na is oxidized to Na+ : Na is

Cl2 is reduced to Cl- : Cl2 is .

2)What are redox reactions?

Activity 2

Mark with arrows the oxidation and reduction processes as shown in the example

below:

CuO + C Cu + CO2

ZnO + Mg ( MgO + Zn

Fe2O3 + 3CO ( 2Fe + 3CO2

2HI + Cl2 ( I2 + 2HCl

2NH3 + 3Br2 ( N2 + 6HBr

Activity 3

(a)Fill in the blanks with suitable words.

(i)The substance that causes oxidation is called the . agent.

(ii)The substance that causes reduction is called is the agent.

(b)(i)2Zn + O2 ( 2ZnO.

Oxidising agent:..

Reducing agent:..

(ii)2Mg + CO2 ( 2MgO + C

Oxidising agent:..

Reducing agent:..

Activity 4

Example :

a)Determine the oxidation number for the underlined elements

(i) H3PO4

(13) + P + (-24) = 0

3 + P + (-8) = 0

P = 8 3

P = +5

( the oxidation number for phosporus is +5

(ii)N2H4

2N + (14) = 0

2N + 4 = 0

2N = -4

N = -2

( the oxidation number for nitrogen is -2

1.State the changes in oxidation nmber for the bolded element in the boxes provided and state whether the element undergoes oxidation or reduction process .

a)Cr2O7 2- ( Cr3+

Process

b)4HCl + MnO2 ( MnCl2 + Cl2 + 2H2O

Process

c)2Fe + 3Cl2 ( 2FeCl3

Process

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

Process

2. Calculate the oxidation number for chlorine and nitrogen in their compounds below.

Chlorine compound

HCl

HClO

HClO2

ClO2

HClO3

HClO4

Nitrogen Compound

NH3

N2O

NO

NO2-

NO2

NO3-

3.Determine the oxidation number for the underlined elements in the table below and name the compound using the IUPAC nomenclature.

Formula of compound

Oxidation number

Name of compound

CuSO4

Cu2O

CrCl3

Cr2 O72-

NO3-

NH4+

MnO2

Activity 5

For each of the reactions below,

- write the half-equations,

- identify the following:

(i) oxidised substance(iii) oxidising agent

(ii) reduced substance(iv) reducing agent

Example :

Zn + 2Ag+ ( Zn2+ + 2Ag

Oxidation Half-equation : Zn ( Zn2+ + 2e-

Reduction Half-equation : 2Ag+ + 2e- ( 2Ag

Oxidised substance : Zinc (Zn)

Reduced substance : Silver ion (Ag+)

Oxidising agent : Silver ion (Ag+)

Reducing agent : Zinc (Zn)

a)

Cl2 + 2I( ( 2Cl( + I2

b)

c)

d)

Activity 6

Changing Iron(II) ions, Fe2+ to Iron(III) ions, Fe3+

Procedure :

1.The test tube contains 2.0 cm3 of ............................................. solution.

2. Add .................................. drop by drop into the test tube and heat the mixture.

3.Record the .............................................................

4. Observation :

The colour of iron(II) sulphate solution change from ....................... to ................. The bromine water change from ....................... to ........................ In order to detect iron(III) ion in the solution, add ...................................................................................in excess. .................................... precipitate of iron(III) hydroxide is formed.

5. Concept :

a) Iron(II) ion is ..................... to iron(III) ion by ................................... At the same time, .............................................. is reduced to ............................... ion.

b) Half equation

Fe2+ ( ............. + ............

Br2 + ............ ( .....................

c) Iron(II) ions .............. electrons to become iron(III) ions. Iron(II) ions

are .........................................

d) Bromine molecules ............. electrons to form bromide ions. Bromine

molecules are ....................................

e)Reducing agent : ...................................

Oxidising agent : ...................................

Changing Iron(III) ions, Fe3+ to Iron(II) ions, Fe2+

Procedure :

1. The test tube contains 2.0 cm3 of ............................................. solution.

2. Add ......................................... into the test tube and heat the mixture.

3. Record the .............................................................

4. Observation :

a. The colour of iron(III) chloride solution change from (i) .......................

to (ii) .................

b. In order to detect iron(II) ion in the solution, add (i)........................................

in excess. A (ii).................................... precipitate of iron(II) hydroxide is

formed.

5. Concept :

a. Iron(III) ion is (i)..................... to iron(II) ion by (ii)................................... At the

same time, zinc is (iii)......................... to (iv)............................... ion.

b. Half equations :

Fe3+ + ....................( ..............

Zn ( .................. + ..................

c. Iron(III) ions (i) .............. electrons to become iron(II) ions. Iron(III) ions

are (ii) .......................

d. Zinc atom (i).................. electrons to form (ii)................ ions. Zinc atoms

are (iii)........................

e. Reducing agent : (i) .................................

Oxidising agent : (ii) ...................................

Activity 7

Displacement of metal from its salt solution

M ( Mn+ + ne

1.The (i) . (more/ less) electropositive element, is oxidized more

(ii) . (easily / harder )and acts as a (iii) ..

( stronger / weaker) reducing agent.

2. The element that is located higher in the electrochemical series can displace other elements that are (i) .. in the electrochemical series from its salt solution.

Example :

Zn + CuSO4 ( ZnSO4 + Cu

Zn ( Zn2+ + 2e [ Oxidation ]

Cu2+ + 2e ( Cu [ Reduction ]

Zinc, Zn is more electropositive than copper,Cu.

Thus Zn atom releases two electrons to form ion Zn2+.

The electrons are transferred from atom Zn to the copper(II) ion, Cu2+.

The copper(II) ion, Cu2+ receives the two electrons to form atom copper, Cu.

Zn atom acts as (ii)

Cu2+ ion acts as (iii)

Activity 8

Displacement of halogens from their halide solutions by other halogens.

lessreceive

halide

seventeendecreases

oxidising agentmore

1)Halogens are located in Group .. of the Periodic Table

2) Halogen elements tend to electrons to achieve a stable octet electron arrangement.

3) Halogens are reduced to ions

4) The electronegativity of halogens or their tendency to accept electrons .. when going down Group 17.

5) Thus, the reactivity of halogens acting as .. decreases when going down the group 17.

6) In a displacement of halogen, a .. electronegative halogen displaces a electronegative halogen from its halide solution.

Activity 9

Transfer of electrons at a distance

1.Redox reaction involving bromine water and potassium iodide solution

2.The half equation for the reaction that occurs around the carbon rod on the left is as follow.

MnO4- + 8H + + 5e ( Mn 2+ + 4H2O

a)State the change in oxidation number of manganese.

..

b) A brown vapour was observed around the carbon rod on the right. Name the brown vapour formed.

..

c) Write a half equation for the formation of the violet vapour.

..

d)On the diagram above, show the direction of the flow of electron by

using arrows.

B.RUSTING AS A REDOX REACTION

Learning Outcomes :


state the conditions for the rusting of iron

.state what corrosion of metal is.

describe the process of rusting in terms of oxidation and reduction.

Generate ideas on the use of other metals to control rusting,

Explain with examples on the use of a more electropositive metal to control metal corrosion,

Activity 10

1. Fill in the blanks with suitable terms

a) (i).. is the oxidation of a metal while (ii) is the oxidation of iron when the metal or iron interact with the environment by losing electron.

b) The presence of (i). and (ii).will cause iron to rust, the rusting can be accelerated by adding (iii). or (iv) .

c) Copper forms a green coating as a result of corrosion, it contains .. ions.

d) Metals like (i). and (ii) forms a very tough oxide which adhere tightly to the surface of the metal preventing further oxidation, thus protecting the metals from further corrosion.

e) To prevent corrosion, metal A can be coated with a layer of metal B which is more (i). than A. Metal B will corrode first, thus preventing metal A from corrosion. In this situation, metal B is also called the (ii). metal.

f) If iron is in contact with another (i)which is less electropositive than iron, the rate of rusting for iron will be (ii).

g) Zinc is used to protect iron or steel by coating a thin layer of zinc onto it, this process is called (i).. In industry, zinc is chosen to serve the purpose rather than other metals because zinc is (ii) in cost.

h) Some household and bathroom equipments are coated with a layer of shiny finishes, the metals usually used for these coatings are (i). and (ii). .

2.Write the equations for the processes below

a) The formation of iron(II) ions from the metal:

When iron contacts with water, the iron surface oxidizes to form iron(II) ions.

Equation: .

b) The formation of hydroxide ions:

Electrons travel to the edges of the water droplets, where there is high concentration of dissolved oxygen. Water and oxygen molecules receive the electrons, they are reduced to form hydroxide ions.

Equation: .

c) Fe2+ readily combines with OH- to form Fe(OH)2

Equation:

d) With excess oxygen, the rust is formed:

The Fe2+ ions are further oxidized to form Fe3+ ions, which reacts with OH- ions to form the hydrated iron (III) oxide, Fe2O3. xH2O, known as rust.

Equation: .

Activity 11

Answer the questions below.

1.Name three main ways to prevent rusting.

(i)

(ii)

(iii)

2.Galvanising involves coating an iron or steel sheet with a thin layer of..

3.Name three metals that can be used in sacrificial protection for an underground pipe.

4.Name a reagent that is usually used to detect the presence of iron (II) ion in an experiment to investigate the rusting of iron.

..

Activity 12

1

Diagram below shows the apparatus set-up for the experiment to study the effect of metals P and Q on the rusting of iron nail. The results are recorded after one day.

Experiment

After 1 day

Observation

A

Some dark blue precipitate.

B

Large amount of dark blue precipitate

C

No dark blue precipitate. Solution turns pink.

(a)

Write the half-equation for the formation of iron(II) ion from iron.

..............................................................................................................................

(b)

State the function of potassium hexacyanoferrate(III) solution in the experiment.

..............................................................................................................................

(c)

Which test tube shows the highest rate of rusting of iron? Explain your answer.

..............................................................................................................................

..............................................................................................................................

(d)

Arrange the metals Fe, P and Q in order of decreasing electropositivity.

..............................................................................................................................

(e)

(i)

What happens to metal Q in test tube C?

.................................................................................................................

(ii)

Suggest a metal that can be used as metal Q.

.................................................................................................................

(f)

State the ion that causes the solution in test tube C to turn pink.

..............................................................................................................................

C.THE REACTIVITY SERIES OF METALS WITH OXYGEN AND ITS APPLICATION

Learning Outcomes :


compare the differences in the vigour of the reactions of some metals with oxygen.

deduce the reactivity series of metals.

determine the position of carbon and hydrogen in the reactivity series of metals.

state what the reactivity series of metals are.

describe the extraction of iron and tin from their ores.

explain the use of carbon as the main reducing agent in metal extraction.

use the reactivity series of metals to predict possible reactions involving metals

Activity 13

1.

Figure below shows the set-up of apparatus for an experiment to determine the order of metals in the reactivity series. Solid potassium manganate (VII) is heated to release oxygen gas to react with metal powder.

The experiment is carried out using metal powders of copper, zinc, magnesium and lead . The observation of the experiments on the metal powders of copper, zinc, magnesium and lead in the experiments are shown in table below.

Type of metal

Observation

Copper

Faint glow

Zinc

A bright flame spreads slowly

Magnesium

A bright white shiny flame spreads quickly

Lead

Red hot and embers slowly

(a) Based on the observations in table above, arrange copper, zinc ,

magnesium and lead in descending order of reactivity of metal towards

oxygen.

Descending order of reactivity of metal towards oxygen.

(b) Name two other substances that can be used to release oxygen gas.

(i)

(ii)

(c) Write a balanced chemical equations for the reactions below.

(i) Copper + oxygen

(ii) Zinc + oxygen

(iii) Magnesium + oxygen

(iv) Lead + oxygen

.

2) According to the equation below, Carbon reacts with oxygen to

form ..

C ( s ) + O2 (s ) CO2 ( g )

3) Carbon displaces a metal from its metallic oxide. Thus, by

heating a mixture of metal oxide and carbon , the reactivity of carbon can be

determined.

0 Oxidation +4

C ( s ) + 2PbO (s ) CO2 ( g ) + 2Pb(s)

+2 Reduction 0

4) Carbon is to carbon dioxide and lead (II) oxide is reduced to

lead.

5) Carbon acts as a (i).. agent which displaces a metal from its oxides while the metallic oxide is the (ii) .agent.

6) Carbon is not able to displace a . metal from its metallic oxide. Hence, there is no displacement reaction when a mixture of magnesium oxide and carbon is heated.

7) Carbon is positioned in between (i) and (ii) .. in the reactivity series.

8)

Oxides of metals W,X, Y, Z are heated with equal amount of carbon powder in an experiment to compare their reactivity with carbon. The changes observed are recorded in table 2.2

Mixture

Observation

Carbon + oxide of W

No change

Carbon + oxide of X

Dim glow.Grey residue is formed

Carbon + oxide of Y

No change

Carbon + oxide of Z

Bright glow. Brown residue is formed

(a) Based on the observation , classify the metals in groups that are

(i) more reactive than carbon

(ii) less reactive than carbon

(b) Suggest the possible elements for metals X and Z

X : .

Z : .

(c) The reactivity of W and Y can be compared by heating an equal amount

of W powder with oxide of Y in a crucible lid using the same apparatus

set-up.

(i) What is the expected observation if W is more reactive than Y?

Explain.

(ii) Suggest the possible elements for W and Y

W : .

Y : ..

(iii) Write a balanced chemical equation for the reaction between W

(charge of +2) and oxide of Y (charge of +3).

..

(iv) Identify the reducing agent and oxidising agent in the reaction

between W and oxide of Y.

Reducing agent :.

Oxidising agent : .

9.

Complete the reactivity series with oxygen below by writing the name of the elements.

Potassium

Calcium

Aluminium

Iron

Tin

Mercury

Gold

Reactivity

decreases

10.The following shows part of the reactivity series of metals with oxygen.

K

Na

Ca

Mg

Al

Zn

Fe

Sn

Pb

Cu

Hg

Ag

Au

Insert the positions of carbon and hydrogen in the above series by using arrows.

11.Predict what will be observed when,

(a)hydrogen gas is heated with copper (II) oxide.

............

.

(b)a piece of burning magnesium ribbon is dropped into a gas jar filled with carbon dioxide.

...

(c)carbon is heated with magnesium oxide.

....................

12.Complete the following table by giving the main mineral ore and the metal extracted from ore.

Ore

Main mineral in ore

Metal extracted

Name

Formula

(a) Bauxite

(b) Hematite

(c) Magnetite

(d) Cassiterite

D.ELECTROLYTIC AND CHEMICAL CELLS

Learning Outcomes:


explain with examples the oxidation and reduction reactions at the electrodes of various chemical cells.

explain with examples the oxidation and reduction reactions at the electrodes of various electrolytic cells.

state the differences between electrolytic and chemical cells in terms of basic structure, energy conversion and the transfer of electrons at the electrodes .

compare and contrast electrolytic and chemical cells with reference to the oxidation and reduction process.

Activity 14

1)The differences between electrolytic and chemical cells

2)The figure above shows the electrolysis process of molten lead (II) bromide. Answer the questions below.

(a)State the ions contain in the molten lead (II) bromide.

.

(b)Pb2+ ions move to the (i) . while Br( ions move to the (ii) ..

(c)Br( ions act as the (i) . agent, losing electrons to become (ii) .. molecules. Thus, Br( ions undergo (iii) .. process.

(d)Oxidation half-equation : .

(e)Pb2+ ions act as the (i) . agent, accepting electrons to become (ii) .. metal. Thus, Pb2+ ions undergo (iii) .. process.

(f)Reduction half-equation : ...

3) Figure below shows a chemical cell.

a) Label the negative terminal and positive terminal and show the direction of the flow of electrons by using arrows the above figure.

b) Write the half-equation for the reaction at the positive terminal.

...

c) Write the half-equation for the reaction at the negative terminal.

...............

d)State the substance that undergoes oxidation process.

...

e)State the substance that undergoes reduction process.

...

4)Similarities and differences of the redox reactions in electrolytic cell an chemical cell

Similarities

In both cells,

electrons are transferred from the agent to the agent.

oxidation occurs at the anode.

occurs at the cathode

Differences

Activity 15

1.Draw and label one example of primary cell. Discuss the oxidation and reduction processes that occur in the cell.

2. Draw and label one example of secondary cell . Discuss the oxidation and reduction processes that occur in the cell.

Reduction

Oxidation

Oxidised Substance:

Reduced Substance:

Oxidising Agent:

Reducing Agent:

Oxidation Half-equation:

Reduction Half-equation:



Reducing Agent:

Oxidising Agent:

Reduced Substance:

Oxidised Substance:

Mg + 2HCl ( MgCl2 + H2



Reducing Agent:

Oxidising Agent:

Reduced Substance:

Oxidised Substance:

Pb + Br2 ( PbBr2



Reducing Agent:

Oxidising Agent:

Reduced Substance:

Oxidised Substance:

4Na + O2 ( 2Na2O

(1) 2.0 cm3 of Iron(II) sulphate solution (light green)

(1) 2.0 cm3 of iron(III) chloride solution (yellow)

1. Electrons flow from ..

to

2. The colour of potassium iodide solution

changes from to

.

3. The colour of bromine water changes

from to

4. Oxidation half-equation :

.

Reduction half-equation :

.

Overall Ionic Equation :

.

6. Oxidising

agent :

7. Reducing

agent :

Carbon

Electrode

Bromine potassium

Water iodide

Solution

Dilute

Sulphuric acid

The electrodes may be of

the . material such as

as

It requires a source of

.

It does not require a source of

The electrodes must be of

two ...... metals.

Potassium Iodide Solution, 1.0 mol dm-3

Acidified Potassium Manganate (VII) solution, 1.0 mol dm-3

Dilute Sulphuric acid,

1.0 mol dm-3

Carbon rod

The chemical reactions that occur at

the electrodes produce

an

The electrical energy

causes reactions

to occur at electrodes.

Electrons flow from the .

electrode (anode) to the .

electrode (cathode).

Electrons flow from the more . metal ( terminal) to the less metal (. terminal)

Chemical cell

Electrolytic cell

Y

X

PAGE

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3 Oxidation and Reduction