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CHAPTER 3: OXIDATION AND REDUCTION B. 3.0

OXIDATION AND REDUCTION

Redox reactions

Rusting as a redox reaction

Activation Series of Metal and its application

Electrolysis cell and chemical cell

3.1 3.2

3.3 3.4

1. Define redox reaction. 2. Stated the difference between redox and non redox reaction.. 3. List 5 examples of redox reaction and non redox respectively.

1. State two conditions to cause the metal rusting. 2. Explain the rusting process of iron. 3. Explain why iron plated by aluminium less rusted compared to the iron plated with copper.. 4. Define metal corrosion

1. State the difference between electrolysis cell and chemical cell. 2. Draw an example of electrolysis cell and then explain the process occurs in the cell. 3. Draw an example of chemical cell and then explain the process occurs in the cell.

1. Define Activation Series of Metal. 2. Draw a diagram showing arrangement of apparatus for experiment to get metal activation series. 3. Draw a diagram and then describe an experiment to locate hydrogen in the activation series of metal.

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B. 3.1

REDOX REACTION (oxidation & reduction)

Oxidation and reduction in terms of oxygen.

Oxidation and reduction in terms of electron transfer

Oxidation and reduction in terms of oxidation numbers

Definition for redox reaction The reaction involved oxidation and reduction simultaneously *3.1.1

Oxidation and reduction refers to hydrogen.

*3.1.2

*3.1.3

*3.1.4 Examples of Redox reaction

*3.1.5

1. Define oxidation process in term of oxygen transfer. 2. Define reduction process in term of oxygen transfer. 3. Give an example of oxidation reaction and then write the chemical equation.

1. Define oxidation process in term of hydrogen transfer. 2. Define reduction process in term of hydrogen transfer. 3. Give an example of oxidation reaction and then write the chemical equation.

1. Define oxidation process in term of electron transfers. 2. Define reduction process in term of electron transfers. 3. Give an example of oxidation reaction and then write the chemical equation.

1. Define oxidation process in term of oxidation number. 2. Define reduction process in term of oxidation number. 3. Define oxidation process in term of oxidation number. 4. Give an example of reduction reaction. Write the chemical equation

1 List 3 example of redox reaction. Write chemical equation for each example. 2. List 3 example of non-redox reaction. Write chemical equation for each example.

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*B. 3.1.1

OXIDATION AND REDUCTION IN TERMS OF OXYGEN.

Experiment for oxidation of metal by oxygen.

Experiment for reduction of metal oxide by carbon

Diagram

Definition for oxidation:

Diagram

Definition for reduction:

Equation

Equation

2CuO + C Cu + CO2

Other example to reduction

Other example to oxidation

C + O2 CO2

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*B. 3.1.2

OXIDATION AND REDUCTION IN TERMS OF HYDROGEN.

Definition for oxidation

Example of oxidation

Definition for reduction

Example of reduction

3CuO + 2NH3 3Cu + N2 + 3H2O 2NH3 + 3Br2 N2 + 6HBr H2S + Cl2 S + 2HCl

3CuO + 2NH3 3Cu + N2 + 3H2O 2NH3 + 3Br2 N2 + 6HBr H2S + Cl2 S + 2HCl

Determine; Oxidizing agent Reduction agent Elements that have been oxidize. Elements that have been reduced

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*B. 3.1.3 G

OXIDATION AND REDUCTION IN TERMS OF ELECTRON TRANSFER

Experiment to observe redox refers to the electron transfer in a distance.

Definition for oxidation

Other example

Definition for reduction

Diagram

Equation

Sulphuric acid

Chlorine water Potassium iodide

Carbon Carbon

2I- I2 + 2e (oxidation)

Cl2 + 2e 2Cl- (reduction)

Overall equation 2I- + Cl2 I2 + 2Cl- (Redox)

e

Electron transfer in a distant

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*B. 3.1.4

OXIDATION AND REDUCTION IN TERMS OF OXIDATION NUMBERS.

Oxidation number

Calculation of oxidation number 1. 2.

The relation of oxidation numbers and the IUPAC nomenclature.

Determination of oxidation number of an element.

Definition for oxidation number

General guidance to determine oxidation number. 1. Oxidation no. of an atom/molecule of an element = 0 2. Oxidation no. of an element in monatomic ion = number of charge at the ion. 3. Oxidation number for halogen is -1 except when reacted with more electronegative element (i.e. NBr3

and Cl2O7 is +1 and +7 respectively) 4. Oxidation no. in oxygen is -2 except in F2O and H2O2 is +2 and –1 respectively 5. Oxidation number of hydrogen is +1 except in metal hydrides such as NaH and MgH2 is equal -1 6. Total of oxidation no. in all elements in a neutral compound = 0 7. Total of oxidation no. of all element in a complex ion = number of charge at the ion

The change of oxidation numbers in the oxidation and reduction process.

Oxidation:

Reduction:

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*B. 3.1.5

THE EXAMPLES OF REDOX REACTION

Conversion of iron (II) to iron (III) ion and vice versa

Displacement reaction of a metal from the solution of their compound

Electrons transfer in a certain distance.

Displacement reaction of the halogen from their halide solution

Experiment

Half equation (oxidation): Half equation (reduction): Overall equation:

Experiment: Displacement of copper by magnesium

Half equation (oxidation): Half equation (reduction): Overall equation:

Sulphuric acid

Oxidation agent

Reduction agent

Carbon Carbon

e

Electron transfer in a certain distance

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Diagram Examples

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B.3.2

RUSTING IS REDOX REACTION.

Definition of metal corrosion

Rusting process of iron refers to oxidation and reaction.

Using of the other metals to protect the iron from rusting. (Example, iron plated by zinc): Zinc is more electropositive than iron, suppose to be the negative terminal. Zinc eliminated electrons and corroded to prevent iron from corroded. Therefore iron not corroded and not rusted. Zinc is called Zink dianggap sebagai logam korban.

Experiment: The effect of contact by other metal to the rusting of iron.

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B. 3.3

Application of the Reactivity series of metals in the metal extraction 1. Extraction of iron 2. 3.

Def. of metal reactivity series

THE REACTIVITY SERIES OF METALS AND ITS APPLICATIONS

Experiment: To get the reactivity series of metals.

Experiment: To determine the position of carbon in metal activation series.

Experiment: To determine the position of hydrogen in the reactivity series. Extraction of iron and tin

in industry

KMnO4

Glass wool

Metal powder

Heat

Heat

Metal oxide + Carbon

CO2

Metal oxide

Heat Dry H2

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B. 3.4

ELECTROLYTIC CELL AND CHEMICAL CELL

Reactions in electrolytic cell

Cu2+ , NO3- ,

H+ , OH-

At cathode: Cu2+ + 2e Cu (Reduction) At anode: 4OH- 2H2O + O2 + 4e (Oxidation)

Carbon

Copper nitrate (electrolyte)

Other examples

Anode ……Oxidation took place Cathode ……Reduction took place

Reaction in a chemical cell

Simple cell

Metal Pb Metal Mg

Electrolyte H2SO4

At terminal (-) Mg Mg2+ + 2e (Oxidation)

At terminal (+) 2H+ + 2e H2 (Reduction)

anode

A

Other examples

Oxidation is at the anode. Reduction is at cathode.

Oxidation is at cathode. Reduction at anode

Comparison between electrolytic and chemical cells refers to the oxidation and reduction.

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3.1 Analysing redox reactions 1. Determine which substance is oxidized or reduced in each of the following reactions. ( In terms of loss or gain of oxygen and hydrogen ) a. Mg + CuO → MgO + Cu b. Zn + PbO → ZnO + Pb c. C + 2 ZnO → CO2 + Zn d. H2S + Cl2 → S + HCl e. 2NH3 + 3 Br2 → N2 + 6 HBr 2. Fill in the blanks. 2Fe(S) + 3Cl2(g) → 2FeCl3

( a)…………………………is oxidized.

…………………………is reduced. …………………………is the oxidising agent. …………………………is the reducing agent.

Lost electron (oxidation)

receive electron (reduction)

Oxidation : Combination of a substance with oxygen. Loss of hydrogen Loss of electrons Reduction : Removal of oxygen from a substance Gain of hydrogen Gain of electrons.

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2Ag+ + Cu → 2Ag + Cu2+

b………………………..is oxidized ………………………..is reduced. …………………………is the oxidizing agent. …………………………..is the reducing agent.

3. Determine which elements is oxidised and reduced in each of the following reactions.(In term of electron transfer).

a. Mg + 2HCl → MgCl2 + H2 b. 2Fe + 3I2 → 2FeI3

c Zn + Cu2+ → Zn2+ + Cu Oxidation : O:xidation number of the element increases Reduction : Oxidation number of the element decreases 4. Fill in the blancks. 2Fe(S) + 3Cl2(g) → 2FeCl3

a……………………….is oxisided ………………………….is reduced

Receive electron (reduction)

Lost electron (oxidation)

Oxidation number increase (Oxidation) from 0 to +3

Oxidation number decrease (reduction) from 0 to -1

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2Ag+ + Cu → 2Ag + Cu2+

b.………………………………………is oxidized. .……………………………………….is reduced. 5. The oxidation number of an atom or molecule in its element is zero. a. Fill in the table for the oxidation number of the atom and molecule.

Atom Oxidation number Molecule Oxidation number Mg 0 H2 0 Cu O2 Na I2 He Br2 Fe Cl2 F2 N2

Oxidation number of hydrogen in a compound is always +1 ( except in metal hydrides where is is -1 ) Oxidation number of oxygen in a compound is always -2 ( except in peroxides) The sum of the oxidation numbers of all the elements in the formula of a compound must be zero. The sum of the oxidation numbers of all the elements in a polyatomic ion must be equal to the charge of the ion.

Oxidation number decrease (reduction) from +1 to 0

Oxidation number increase (oxidation) from 0 to +2

Oxidation number is the charge that the atom of the element would have if complete transfer of electrons occur.

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b. Calculate the oxidation number,ON of the following. (i ) Mangan in MnO4

- Let ON mangan is x 1(x) + 4(-2) = -1 x = =+ 7 (ii) Sulphur in S2O3

2- (iii) Mangan in MnO2

(iv) Mangan in Mn2O3

(v) Nitrogen in NH3 6. Name the following compounds.The first two has been done .

Formula Oxidation number,ON of element

Name

FeSO4

Fe : +2 Iron (II) sulphate

NaClO

Na : +1 O : -2 Let ON of Cl = x (+1) + x +(-2) = 0 x = +1

Sodium chlorate (I)

The Roman numerals represent the oxidation numbers of Elements.The oxidation number is included in the IUPAC nomenclature of a compound only if the element in involved has more than one oxidation number

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Formula Oxidation number,ON of element Name

PbO2

PbO

MnO2

K2Cr2O7

7. 7.1 Change of Fe2+ to Fe3+

Fill in the blanks.

1. When bromine water is added to iron(II)sulphate solution, the bromine water turns from………………….to …………………………………Iron(II)sulphate solution turns from ……………………..to……………………….. 2. The iron(II)ions is oxidized to……………………………….by…………………..

electrons.

3. Bromine water is reduced as bromine molecules …………………….electrons

to form bromide ions.

Redox reactions are chemical reactions involving oxidation and reduction occurring simultaneously. 7.1 Change of Fe2+ to Fe3+ / Fe3+ to Fe2+ 7.2 Displacement of metal from its salt solution 7.3 Displacement of halogen from its halide solution 7.4 Transfer of electrons at a distance.

Iron(II)sulphate solution

Bromine water

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4. Half-equation for oxidation : Fe2+ (aq) → Fe3+(aq) …………………..

5. Half –equation for reduction : Br2(aq) …………… → 2 Br- (aq)

6. Overall equation for redox reaction :

……..Fe2+ (aq) + Br2 (aq) → 2 Fe3+ (aq) + …… Br- (aq)

Change of Fe3+ to Fe2+

Complete the table for the reactions that take place in the test tube above.

Observations

Half-equation : oxidation

Half –equation : reduction

Overall equation

Comfirmatory test

Zinc powder

Iron(III)chloride solution

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7.2 Displacement of metal from its salt solution.

For each of the following reactions, write down the half –equations and overall equation.State which substance is oxidised and reduced.

a. Magnesium is added to aqueous copper(II) sulphate.

Oxidation Reduction Half-equation

Observations

Overall equation

K Na Ca Mg Al Zn Fe S Pb H Cu Hg Pt Ag Increasing electropositivity. A more electropositive metal displaces a less electropositive metal from its salt solution.

Zinc is more electropositive than copper ( zinc is in higher position than copper in electrochemical series). Zinc displaced copper from copper(II) sulphate solution. Zn(s) → Zn2+(aq) + 2e Zinc is oxidised

The copper(II) ions are taken out from the solution to form copper metal . Cu2+ (aq) + 2e → Cu (s) Copper(II) ions are reduced. The blue colour of copper(II)sulphate solution fades.

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b. Copper is added to silver nitrate solution. Oxidation Reduction Half-equation

Observations

Overall equation

c. Zinc is added to copper (II) sulphate solution. Oxidation Reduction Half-equation

Observations

Overall equation

7.3 Displacement of halogen from its halide solution

A more reactive halogen displaces a less reactive halogen from its aqueous halide solution.The more reactive halogen has a higher tendency to gain electron. F2 Highest tendency to gain electrons Cl2 Br2 I2 At2

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Cl2 + 2 KI → 2 KCl + I2

Chlorine is more reactive than iodine.Chlorine molecules Cl2(aq) + 2e → 2 Cl- (aq) receive electrons to form chloride ions.Chlorine is reduced.Chlorine is the the oxidising agent. Iodide ions in potassium bromide 2I- (aq) → I2 (aq) + 2e lose electrons to form iodine molecules.Potassium iodide solution is the reducing agent. a. Bromine water id added to potassium iodide solution.

Oxidation Reduction Half-equation

Observation

Overall reaction

Reducing agent : Oxidising agent :

Potassium iodide solution

Chlorine water

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b) Chlorine water is added to potassium bromide solution.

Oxidation Reduction Half-equation

Observation

Overall reaction

Reducing agent : Oxidising agent :

2.Halogens give different colours in 1,1,1,-trichloroethane. Complete the table below. Halogen Colour of halogen in

aqueous solution Colour of halogen in 1,1,1- trichloroethane.

Chlorine

Bromine

Iodine

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7.4 Transfer of electrons at a distance

1. Classify the following substances into oxidising agent and reducing agent.

Oxidising agent

Reducing agent

e e -

When an oxidising agent and a reducing agent are kept ‘ at a distance’, the electrons have to be transferred from one to the other through an external circuit. A redox cell is a device to produce electric current from a redox reaction

Bromine water potassium iodide solution iron(II) sulphate solution Potassium manganate (VII) solution Tin(IV)chloride

G Acidified potassium (Oxidising manganate(VII),KMnO4 agent ) Reduction : MnO4

- (aq) + 8H+(aq) + 5e → Mn2+(aq) + 4 H2O(l) ON of mangan is reduced from +7 to +2. The purple acidified potassium manganate(VII) solution decolourises.

Iron(II)sulphate ( reducing agent) Oxidation : Fe2+

→ Fe3++e ON of iron is oxidised from +2 to +3.

Carbon electrodes

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Complete the following table for the transfer of electrons at a distance.

Negative Terminal

Positive Terminal

Reducing agent

Oxidising agent

Half-equation

Name of the products.

Observations

Comfirmatory test

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Potassium iodide (aq) Bromine water

Sulphuric acid

Carbon electrodes

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Complete the following table for the transfer of electrons at a distance.

Negative Terminal

Positive Terminal

Reducing agent

Oxidising agent

Half-equation

Name of the products.

Observations

Comfirmatory test

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Carbon electrodes

K2Cr2O7 / H+

Ferum(II) sulphate

Potassium chloride,KCl

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Complete the following table for the transfer of electrons at a distance.

Negative Terminal

Positive Terminal

Reducing agent

Oxidising agent

Half-equation

Name of the products.

Observations

Comfirmatory test

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Copper(II) sulphate

Copper

Magnesium sulphate

Magnesium

Potassium iodide

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3.2 Analysing rusting as a redox reaction. 1.Fill in the blancks.

Negative pole : Iron atoms in contact with the centre of the water droplets ionsise to form……………..Oxidation occurs. .The half- equation is Fe(s) → ……………+ 2e Electrons are transferred from the iron atoms to oxygen and water molecules at the edge of the water droplets. Reduction occurs when the surface of the water droplets exposed to the air has a tendency to……………………..electrons.The half-equation is : O2(g) + 2H2O (l) + …………… → 4 OH- (aq). The ………………….ions from the ionisation of iron then combined with the hydroxide ions to form…………………………………….as a dirty green precipitate.The overall equation for the redox reaction is : 2Fe(aq) + O2(g) +2H2O → 2 Fe(OH) 2 (s) ……………………is then oxidised in the air to form……………………………….. which then becomes rust,……………………………….,Fe2O3.3H2O.

Rusting is the corrosion of iron.Rust forms on the surface of iron. Fe(s) → Fe2+(s) + 2e It involves electron loss which is oxidation.

Negative pole

Positive pole

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In order to rust, both air and water a must Air alone wont do, without water there too, So protect it,or get a brown crust! When iron and copper are in contact with each other in the presence of electrolyte,electrons are transferred from iron to copper.Rusting of iron occurs.

All metal atoms ionise to form metallic ions .The higher the metal in in the electrochemical series, the easier its atoms ionise and the easier the corrosion occurs. Rusting is prevented if iron is in contact with a more electropositive metal. Rusting is faster if iron is in contact with a less electropositive metal.

Fe(s) → Fe2+(s) + 2e Iron loses electrons more readily than copper. Fe2+ ion formed in the electrolyte react with potassium hexacyanoferrate(III) solution to form dark blue spots.

Hot agar solution which contains potassium hexacyanoferrate(II)

copper

Dark blue spots Blue spots Pink colour observed. Less intensity Gas bubbles are formed. of pink colour

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2. Complete the table.( The observations are given in the box.)

Pair of Metals

Is the iron corrode?

Observations

Mg/Fe

Zn/Fe

Sn/Fe

Pb/Fe

3. .

Alloys

Painting

Clothes Hanger

Roofs of houses

Medical instruments

Method of preventing rusting of iron

A protective layer A sacrificial metal

Oiling

Tin-plating

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Comparison between electrolytic and voltaic cells Similarities: Electrolytic cell

Voltaic cell

� Contains an electrolyte � Consists of an anode and a cathode � Positive ions and negative iond move in the electrolyte � Chemical reactions involve the release and acceptance of electrons

Differences:

Characteristics Electrolytic cell Voltaic cell Energy change

Electric current and reactions

Electric current results in a chemical reaction

Chemical reaction produces an electric current

Cathode and anode Cathode: Anode:

Cathode: Anode:

Flow of electrons

Negative terminal

Positive terminal

Types of electrodes

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EXECISES

1. Below is the half equation of a reaction What is meant by oxidation reaction based on the equation ?

A. Electrons are received by bromine B. Electrons are donated by bromine C. Electrons are received by bromide ions D. Electrons are donated by bromide ions.

2. Which of the following are oxidizing agents ?

I. zinc II. Bromine water III. Potassium iodide solution IV. Acidified potassium manganate(VII)solution.

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

3. Below is an ionic equation

Which of the following is true of the equation ?

A. Y2+ is oxidized B. X is an oxidizing agent C. X2+ is a reducing agent D. X donates electrons to Y2+

4. Fe3+ ions in solution can be converted to Fe2+ ions by adding zinc powder. Which of the following can replace zinc powder in this reaction ? A. Bromine water B. Potassium iodide solution C. Potassium hexacyanoferrate(II)solution D. Acidified potassium manganate(VII) solution

2 Br- → Br2 + 2e

X (s) + Y2+ (aq) → X2+ (aq) + Y (s)

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5. Table 1 shows the result of an experiment for three chemical cells.

Chemical cell Metal pairs Negative terminal Cell voltage/V X P and R R 1.9 Y R and S S 0.8 Z Q and R R 0.3

Which of the following can be deduced from table 1 ?

I. The cell voltage is 1.6V when P and Q are used as electrodes. II. The cell voltage is 1.1 V when P and S are used as electrode. III. Electrons flow from terminal Q to terminal S in the metal pair Q and S. IV. P functions as a positive terminal when it is paired with Q,R or S in a cell.

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

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STRUCTURAL QUESTIONS 1. The figure below shows the set-up of apparatus to investigate the reactions that take place in test tubes P and Q

(a) State the observation for the reaction

(i) In test tube P. [ 1 mark ] (ii) In test tube Q. [ 1 mark ] (b) Write the ionic equation for the reaction in (a)(i). [ 1 mark ] (c ) State what is meant by oxidising agent in terms of electrón transfer. [ 1mark ]

Copper(II) sulphate solution

Magnesium ribbon

Bromine water

Ferum(II) sulphate solution

Test tube P Test tube Q

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(d ) Referring to the reaction that takes place in test tube P. (i) What is the change in the oxidation number of magnesium? [ 1 mark ] (ii) name the oxidizing agent. [ 1 mark ] (e) Referring to the reaction that takes place in test tube Q, (i) State the type of reaction that occurs. [ 1 mark ] (ii) State the oxidation number of bromine in bromine water. [ 1 mark ] (iii) what is the function of bromine water? [ 1 mark ] (iv) name another reagent that can replace bromine water.. [ 1 markah ]

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3. The figure above shows the set-up of apparatus to investigate the electrolysis of dilute copper(II)

sulphate solution. (a) (i) What is meant by cation ? [ 1 mark ] (ii) What is the energy change that occurs in the electrolysis process? [ 1 mark ] (b) In the electrolysis of dilute copper(II) sulphate solution:

(i) State all the ions in the electrolyte. [ 1 mark ]

(ii) In the table below, write the ions in b(i) which moved to electrodes X and Y.

Elektrode X Elektrode Y

[ 1 mark ]

(iii) What are the processes that occur at electrodes X and Y ? Elektrode X : Elektrode Y : [ 1 mark ]

Carbon electrode X

Copper(II) sulphate solution

Carbon electrode Y

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(iv) What would you observe at electrode Y [ 1 mark ] (iv) What is the colour change of the electrolyte? [ 1 mark ]

(c ) (i) Name the gas collected in the test tube at electrode X. [ 1 mark ] (ii) The volume of gas collected at electrode X is 20.0 cm 3, How many moles of gas were collected?

Use the information that 1 mole of gas occupies a volume of 24. 0 dm3 at room temperature and pressure [ 1 mark ]

(iii) Based on the answer in c(ii) what is the number of gas molecules collected? Use the information that the Avogadro number is 6.02 x 10 23 mol -1 . [ 1 mark ]

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3(a) Lime juice was electrolysed using carbon electrodes. What is produced at the cathode? Write a half equation for the reaction. [2 marks]

(b ) The figure below shows two types of cell.

Compare and aontrast cell P and cellQ. Include in your answer the observation and half-equatins for the reactions of the electrodes in both cells. [ 8 marks ]

(c ) A student Intends to electroplate an iron key with a suitable metal to beautify it. Design a laboratory experiment to electroplate the iron key. Your answer should consists of the following:

• Chemicals required. • Procedures of the experiment • Diagram showing the set-up of apparatus. • Chemical equation involved in the reaction • Observation

[ 10 marks ]

Copper plates

Zinc plate

Copper(II) sulphate solution

Copper plate

Cell P Cell Q

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4. (a) You have an iron key that rusts easily.

State how you would solve this problemusing an electrolysis process. [ 4 marks ]

(b) Electrolysis is carried out on a dilute sodium chloride solution using carbon electrodes. Explain how this electrolysis occurs. Use a labeled diagram to explain your answer.

[ 6 marks ] (c ) Aluminium is placed above zinc in the electrochemical series.

Aluminium and zinc can be used to build a chemical cell , using suitable apparatus and the following chemicals;

Aluminium sulphate solution Zink sulphate solution Sulphuric acid solution

Describe how you would build this chemical cell. Include a labeled diagram in your answer. On your diagram , mark the direction of electron flow, the positive terminal and the negative terminal

[ 10 marks ] 5. a) The following are the formulae of two compounds. Al2O3 Cu2O

(i) Based on the two formulae , state the oxidation number for aluminium and copper. [2 marks]

(ii) Name both the compounds based on the IUPAC nomenclature system. [2 marks]

(iii) Expalin the difference between the names of the two compounds based on the IUPAC nomenclature system.

b) The diagram below shows the set up of the apparatus fo an experiment to investigate electron transfer through a solution.

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(i) Name the oxidation agent in the experiment. [1 mark] (ii) Write the half equations for the reactions that occur at the negative and positive terminals. [5 marks]

(iii) Based on your answer in 5b(ii), describe the oxidation and reduction processes in terms of the electron transfer that occurs at the negative and positive terminals.

State also the changes that can be observed after 10 minutes [8 marks]

G

Graphite electrode (Negative electrode)

Graphite electrode (Positive electrode)

Iron(II) sulphate Solution, FeSO4

Acidic potassium manganate (VII) solution, KMnO4

Dilute sulphuric acid, H2SO4