1202 Question Bank Chemistry Form 5 KSSM

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Contents








iii – xii
Must Know
Chapter 1 Redox Equilibrium 1 – 21
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NOTES 1
Paper 1 4
Paper 2 10
Paper 3 20



Chapter 2 Carbon Compound 22 – 46

NOTES 22
Paper 1 26
Paper 2 35
Paper 3 45



Chapter 3 Thermochemistry 47 – 77

NOTES 47
Paper 1 50
Paper 2 62
Paper 3 73



Chapter 4 Polymer 78 – 94

NOTES 78
Paper 1 80
Paper 2 87
Paper 3 93



Chapter 5 Consumer and Industrial Chemistry 95 – 113

NOTES 95
Paper 1 98
Paper 2 104
Paper 3 112


Answers 114 – 132

KNOW Mnemonics







Redox Equilibrium To Determine Oxidising Agents and
Reducing Agents Based on the Value of E 0

0
Molecules or ions with a more positive (less negative) E value:
O Oxidation
AcRO
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I Is
Ac Accept electron
L Loss R Reduction reaction
O Oxidising agent
0
Atoms or ions with a more negative (less positive) E value:
R Reduction
DoOR

I Is Do Donate electron
O Oxidation reaction
G Gain
R Reducing agent
Mnemonics (Chapter 1) 1 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 1) 7 @ Pan Asia Publications Sdn. Bhd.



Voltaic Cell and Electrolytic Cell Prefix for Naming Carbon Compounds

FAT CAT
Meth Monkey
Electrons flow From Anode To CAThode.
Eth Eats
Battery
+ –
V
e – e – e – e –
Prop Pile
+ –
Anode Electrolyte Cathode
But Bananas
–
Electrolyte
Electrolyte
Anion
+
Cation
Monkey Eats a Pile of Bananas
Voltaic cell Electrolytic cell
Mnemonics (Chapter 1) 3 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 2) 9 @ Pan Asia Publications Sdn. Bhd.

Electrolytic Cell Thermochemistry


• Cation = Ca+ion Heat Heat
Cation is a positive ion

• ANIon = A Negative Ion

• Anion (negative ion) moves to Anode (positive electrode)

• Cation (positive ion) moves to Cathode (negative electrode)

EXothermic reaction ENdothermic reaction
• OxidAtion occurs at Anode
Heat EXit from system Heat ENter the system
• ReduCtion occurs at Cathode


Mnemonics (Chapter 1) 5 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 3) 11 @ Pan Asia Publications Sdn. Bhd.

KNOW Important Definitions







Thermoplastic, Thermosetting and Elastomer Types of Chemical Reaction


• Thermoplastics are loose polymer that can be heated and Exothermic reaction:
reshaped repeatedly.
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• Chemical reaction that releases heat energy to the surroundings.
• Thermosettings are rigid polymer that will not return to • Temperature of the surroundings increases.
their original shape when heated.

• Elastomers are polymer that can be deformed and returned Endothermic reaction:
to their original shape. • Chemical reaction that absorbed heat energy from the
surroundings.
• Temperature of the surroundings decreases.



Important Definitions (Chapter 4) 20 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 3) 14 @ Pan Asia Publications Sdn. Bhd.



Soap and Detergent Heat of Displacement

• Soaps are sodium or potassium salts of fatty acids. • Heat of displacement is the heat change when one mole of
a metal is displaced from its aqueous salt solution by a more
• Detergents are sodium salts of sulphonic acids or alkyl electropositive metal.
hydrogen sulphates.
Q
• Heat of displacement = −
• Saponification is the process to prepare soaps by alkaline n
hydrolysis of fats or oils.
Heat change,
Q = mcθ
Fats or oils Fatty acid + Glycerol = V × 4.2 J g °C × θ
−1
−1
Precipitation Number of moles,
Fatty acid salt (soap) MV
n =
1 000


Important Definitions (Chapter 5) 22 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 3) 16 @ Pan Asia Publications Sdn. Bhd.


• Nanoscience is a study on processing of substances at Heat of Combustion
nanoscale, between 1 nanometre - 100 nanometres.

• Nanotechnology is the study and manipulation of matter at • Heat of combustion is the total heat released when one mole of
nanometer scale to produce new materials or devices. fuel is burnt completely in excess oxygen gas.
−1
• Graphene is the carbon allotrope. Graphene is one-layer thick • Heat of combustion is given in the unit kJ mol .
graphite arranged in hexagonal honeycomb-like structure to
Q
form graphene sheets. • Heat of combustion = −
n
• Graphene sheets can used to produce graphites, carbon Heat change,
nanotubes and fullerene balls. Q = mcθ
−1
−1
= V × 4.2 J g °C × θ
Number of moles,
Mass
n =
Molar mass

Important Definitions (Chapter 5) 24 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 3) 18 @ Pan Asia Publications Sdn. Bhd.

KNOW Important Diagrams







Electrolytic Cell The Vulcanisation Process of Natural Rubber



Battery
+ – C C C C
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e – e – C C C C S S
C C
S C
+ – C C
Anode Electrolyte Cathode C C C C C C C
Vulcanisation C C
S
S
– C C
Anion C C C S C
+ C
Anode Cation Cathode C
C C C S
• Attracts anions • Attracts cations C S
• Positive electrode • Negative electrode Natural C C C C
• Oxidation occurs • Reduction occurs rubber Vulcanised rubber
Important Diagrams (Chapter 1) 49 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 4) 55 @ Pan Asia Publications Sdn. Bhd.


Apparatus Set up for Dehydration of Ethanol Preparation of Detergents

Heat the porcelain chips first before heating the glass wool soaked Step 1
in ethanol gently.
Sulphonation
Alkyl benzene
Concentrated
Alkylbenzene + sulphonic acid +
Porcelain chips H SO
2 4 water
Combustion Ethene gas
tube
Step 2
Glass wool
Glass wool Heat
soaked in
soaked in Neutralisation
propanol
ethanol
Water Alkyl Sodium alkyl
benzene benzene
+ NaOH
sulphonic sulphonate +
acid water
Important Diagrams (Chapter 2) 51 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 5) 57 @ Pan Asia Publications Sdn. Bhd.

Energy Level Diagram of Endothermic Reaction Formation of Grease Droplets During
the Cleansing Action of Soap



Energy
Droplet
of grease Hydrophobic
Products
part

ΔH positive
Soap anions
Reactants
Hydrophilic
part



Important Diagrams (Chapter 3) 53 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 5) 59 @ Pan Asia Publications Sdn. Bhd.

1
Chapter Redox Equilibrium




NOTES



1.1 Oxidation and Reduction MnO (aq) + 8H (aq) + 5e → Mn (aq) +
−
−
2+
+
4
4H O(l)
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2
1. Definition of oxidation and reduction. (b) Acidified potassium dichromate(VI) solution:
2−
+
3+
−
Oxidation Reduction Cr O (aq) + 14H (aq) + 6e → 2Cr (aq) +
2
7
7H O(l)
• Gain of oxygen • Loss of oxygen 2 − −
(c) Chlorine water: Cl (aq) + 2e → 2Cl (aq)
2
• Loss of hydrogen • Gain of hydrogen (d) Bromine water: Br (aq) + 2e → 2Br (aq)
−
−
2
• Loss of electron • Gain of electron
6. A reducing agent is a substance that reduces another
• Increase in oxidation • Decrease in substance and oxidises itself.
number oxidation number
Example:
2+
(a) Magnesium: Mg(s) → Mg (aq) + 2e -
2. Oxidation number or oxidation state is the charge of (b) Sulphur dioxide:
the element in a compound if the transfer of electrons SO (aq) + 2H O(l) → SO (aq) + 4H (aq) + 2e −
2−
+
2
2
4
occurs in an atom to form chemical bonds with other (c) Iron(II) ion: Fe (aq) → Fe (aq) + e −
2+
3+
atoms.
7. Examples of redox reactions:
3. Rules for assigning oxidation number: (a) Conversion of Fe ion to Fe ion
2+
3+
(a) The oxidation number of each atom of a free Ionic equation:
element is 0. Br (aq) + 2Fe (aq) → 2Br (aq) + 2Fe (aq)
3+
−
2+
2
(b) The oxidation number of a monoatomic ion is Reduction half-equation:
equal to the charge of the ion. Br (aq) + 2e → 2Br (aq)
−
−
2
(c) Fluorine in its compounds has a fixed oxidation Oxidation half-equation:
number of −1. Fe (aq) → Fe (aq) + e −
2+
3+
(d) Alkali metals (Group 1 elements) in their (b) Displacement of copper by zinc from
compounds have a fixed oxidation number of +1. copper(II) sulphate solution
(e) Alkaline earth metals (Group 2 elements) in Ionic equation:
their compounds have a fixed oxidation number Zn(s) + Cu (aq) → Zn (aq) + Cu(s)
2+
2+
of +2. Reduction half-equation:
(f) Hydrogen in a compound normally has an Cu (aq) + 2e → Cu(s)
−
2+
oxidation number of +1 when it combines with Oxidation half-equation:
non-metals, but hydrogen has an oxidation Zn(s) → Zn (aq) + 2e −
2+
number of −1 when it combines with metals (c) Displacement of bromine by chlorine from
hydrides. potassium bromide solution
(g) Oxygen in a compound normally has an oxidation Ionic equation:
number of −2 except peroxide compounds Cl (aq) + 2Br (aq) → 2Cl (aq) + Br (aq)
−
−
2
(oxidation number is −1) and compounds with Reduction half-equation: 2
fluorine (oxidation number is +2). Cl (aq) + 2e → 2Cl (aq)
−
−
2
(h) Halogens (Cl, Br, I) in their compounds have Oxidation half-equation:
oxidation number of −1 except when combined 2Br (aq) → Br (aq) + 2e −
−
with fluorine and oxygen. 2
(i) The sum of oxidation numbers of all atoms in a 8. Displacement of metal from its salt solution
neutral compound is always 0. (a) A more electropositive metal displaces a less
electropositive metal from its salt solution.
4. A redox reaction is a chemical reaction in which (b) The electropositivity of a metal is determined
oxidation and reduction occur simultaneously. from its position in the electrochemical series.
(c) A metal can displace any metal below it from its
5. An oxidising agent is a substance that oxidises salt solution.
another substance and reduces itself. Example:
(a) Acidified potassium manganate(VII) solution:
1

(d) Three observations may be observed during a 5. The standard cell potential can be obtained based on
displacement reaction of a metal. the following equation.
• The more reactive metal dissolves
• The less reactive metal is deposited E 0 cell = E 0 cathode − E 0 anode
• The colour of the salt solution may change
0
6. The electrode potential value, E is used to predict:
9. Displacement of halogen from its halide solution (a) Substance that undergoes oxidation or reduction.
(a) Halogens are oxidising agents. Conversely, (b) Substance that acts as an oxidising or reducing
halide ions are reducing agents. agent.
(b) A halogen molecule, X , gains electrons to form
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2 (c) Strength of oxidising or reducing agents.
−
−
the halide ion, X : X (aq) + 2e → 2X (aq)
−
2 Example:
(c) Reactivity or oxidising power of the halogens
decreases going down Group 17. F + 2e 2F − E = +2.87 V
0
−
(d) The halide ion gains electrons to form the 2
−
−
halogen molecule, X : 2X (aq) + 2e → X (aq)
0
−
2 2 Oxidising CI + 2e 2CI − E = +1.36 V
(e) Reducing power of the halide ions increases agent 2
moving down Group 17.
Br + 2e 2Br − E = +1.07 V
0
−
2
1.2 Standard Electrode Potential
• E° value increases, strength as oxidising agent
0
1. Standard electrode potential, E is defined as increases.
• Increasing order of oxidising agent strength:
the difference of electrode potential (voltage) of an Br , Cl , F
2 2 2
electrode system consisting of an electrode half-cell
pairing up with the standard hydrogen electrode (SHE)
half-cell. 1.3 Voltaic Cell
Voltmeter
V 1. A simple voltaic cell consists of two different metals
immersed in an electrolyte solution and connected
H (g)
2 Salt bridge with connecting wires.
298 K and 1 atm Electrode X
Platinum electrode 2. A Daniell cell is an example of a voltaic cell where
Acid solution zinc metal and copper metal are used as electrodes
(concentration of Solution of metal X dipped into their respective ionic salt solutions.
H is 1.0 mol dm ) SHE ion 1.0 mol dm –3
+
–3
Voltmeter
e – e –
2. The standard condition to measure the standard
0
electrode potential, E of the cell: Salt bridge
(a) Concentration of ion in aqueous solution is 1.0 Anode (–) Zn Cu (+) Cathode
−3
mol dm . (Oxidation) KCl (Reduction)
(b) Gas pressure of 101 kPa or 1 atm.
(c) Temperature at 298 K or 25 °C.
(d) Platinum is used as an inert electrode.
ZnSO CuSO
4 4
0
3. The E value of standard hydrogen electrode, SHE
2+
0
−
= 0.00 V. Zn + 2e Zn E = −0.76 V
Cu + 2e Cu E = +0.34 V
2+
0
−
1
0
–
H (aq) + e ⇌ H (g) E = 0.00 V
+
2+
2 2 Zn(s) → Zn (aq) + 2e (Oxidation)
−
2+
−
Cu (aq) + 2e → Cu(s) (Reduction)
Overall ionic equation:
4. The above cell can be represented in the form of a cell 2+ 2+
notation. Zn(s) + Cu (aq) → Zn (aq) + Cu(s)
Cell notation of Daniell cell:
n+
Pt(s) | H (g) | H (aq) || X (aq) | X(s)
+
2+
2+
−3
−3
2 Zn(s) | Zn (aq, 1.0 mol dm ) || Cu (aq, 1.0 mol dm ) | Cu(s)
⎧ ⎪ ⎪ ⎨ ⎪ ⎪ ⎩ ⎧ ⎪ ⎨ ⎪ ⎩
• Represents the • Represents electrode Cell voltage, E 0 = E 0 − E 0
SHE X cell 0 cathode 0 anode
= E − E
Cu Zn
= 0.34 − (−0.76)
= +1.10 V
2

• A protective surface prevents water and
Rust (Fe O .xH O)
Drop of water 2 3 2 air from coming into contact with the iron.
Example: paint, grease, oil, plastic, metals
and using other metals.
OH − • Example of using other metals: Galvanisation
O O 2
2 involves coating iron with a layer of zinc. Tin
Fe 2+ plating or tinning is a process of coating iron
e −
with a thin layer of tin. Chrome plating is used
on car bumpers and decorative items.
Anode (–) Cathode (+) (b) Alloying
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2+
Fe(s) Fe (aq) + 2e − O (g) + 2H O(l) + 4e −
2 2 • Alloy such as stainless steel is made of a
−
4OH (aq)
Iron metal mixture of iron with chromium and nickel.
This alloy has a very high resistance to
2. Corroded iron is covered with a layer of red-brown corrosion. Chromium and nickel form oxide
compound. The red-brown compound is hydrated layer that is very strong and impermeable to
iron(III) oxide, Fe O .xH O or rust.
2 3 2 air and water.
3. Corroded copper is covered with a green layer. (c) Sacrificial protection
The green compound is basic copper(II) carbonate, • Involves placing a more electropositive metal
CuCO .
3 in contact with iron.
4. Ways to prevent rusting: • Example: Magnesium acts as a sacrificial
(a) Use a protective surface metal to protect underground iron pipes from
rusting.

PAPER 1


Each question is followed by four options A, B, C or D . Choose the best option for each question.


1.1 Oxidation and Reduction 4. Carbon is a non-metal and it can be used to
SPM reduce metal oxides. In the reactivity series of
SPM
CLONE
metals, carbon is located between
1. Ethanol is the second member of the alcohol
SPM
SPM family. Ethanol reacts with acidified potassium A calcium and aluminium
CLONE B aluminium and zinc
dichromate(VI) to form ethanoic acid. What type C zinc and iron
of reaction is involved in this conversion? D iron and lead
A Esterification C Halogenation
B Dehydration D Oxidation
5. Diagram 1 shows the experimental set-up used
SPM
SPM to investigate a redox reaction. During a redox
2. Oxidising agent is used to oxidise iron(II) sulphate CLONE
SPM
SPM to iron(III) sulphate. Which of the following is reaction, transfer of electron at a distance takes
CLONE place between the oxidising agent and reducing
not an oxidising agent for the reaction? agent which were separated by solution X.
A Chlorine
B Potassium bromide G
C Hydrogen peroxide
D Concentrated nitric acid Graphite electrodes HOTS Analysing
– +
3. What is the oxidation number of hydrogen in
SPM
SPM aluminium hydride, AlH ? Potassium Potassium
CLONE 3 iodide manganate
A −1 C +1 solution (VII) solution
B 0 D +2 Solution X

4 Question 4: Question 5: Diagram 1
SOS TIP Refer to the positions of metals in the reactivity series of metals. Reduction half-equation: 2+ 2
You need to memorise this series.
MnO + 8H + 5e → Mn + 4H O
−
+
−
4
H ions are supplied by aqueous acid solution
+

47. Diagram 16 shows an experiment to investigate A Magnesium forms an alloy with iron
the rate of corrosion of different metals. B Magnesium reacts more readily than iron
C Magnesium prevents oxygen from coming
into contact with iron
Distilled Distilled
water water D Magnesium forms a protective coating of
magnesium oxide on the iron surface
Iron Magnesium
Copper Lead
I II III IV 49. Diagram 18 shows the reactions of metal P.
Diagram 16 HOTS Analysing
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In which test tube will the rate of corrosion be Corrodes slowly
Metal P P oxide
the slowest?
A I Acid Alkali
B II
Colourless solution
C III
D IV Diagram 18
What is metal P?
48. Underground iron pipes are connected to a block A Aluminium C Calcium
of magnesium as shown in Diagram 17. B Iron D Copper


50. A method to protect iron from rusting is to apply
a protective coating as shown in Diagram 19.
Underground HOTS Analysing
iron pipe Coating of
substance X
Magnesium block Iron
Diagram 19
Diagram 17
Which of the following could not be X?
Why does the magnesium block stops the iron A Aluminium C Magnesium
from rusting?
B Gold D Sodium


PAPER 2

Section A

Answer all questions.

1. Diagram 1 shows the electrolysis of two different acids, acid X and acid Y using graphite electrodes. The
concentration of both acids is 1.0 mol dm .
-3



Acid X Acid Y

Graphite Graphite Graphite
electrode electrodes electrode
W X Y Z


Cell I Cell II
Diagram 1
10
SOS TIP

PAPER 3



0
The standard electrode potential, E is determined by measuring the difference of electrode potential value on
an electrode system consisting of a standard hydrogen electrode (SHE) half-cell and an electrode half-cell.
Voltmeter



H (g)
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2
1 atm; 25 °C
Salt bridge



Platinum Electrode
electrode
Acid solution
SHE
Diagram 1
The SHE half-cell in Diagram 1 is paired with different electrode half-cell as shown in Table 1. The voltmeter
diagram that shows the reading is drawn.

Electrode
Anode Cathode Volmeter diagram Volmeter reading (V)
system
2 3
1 4
SHE + Mg Mg SHE
0 5

2 3
1 4
SHE + Ag SHE Ag
0 5
2 3

SHE + Cu SHE Cu 1 4
0 5

2 3
1 4
SHE + V V SHE
0 5
2 3

SHE + Ti Ti SHE 1 4
0 5

Table 1
20 Question 1:
SOS TIP The E value of standard hydrogen electrode, SHE is 0.00 V.
0

3
Chapter Thermochemistry




NOTES



3.1 Heat Change in Reactions
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Exothermic reactions Endothermic reactions

1. Reaction that releases heat energy to the 1. Reaction that absorbs heat energy from the
surroundings. surroundings.
2. Energy content of reactants is higher than that 2. Energy content of products is higher than that
of products. of reactants.
3. ΔH has a negative value. 3. ΔH has a positive value.
4. Heat released to the surroundings, the 4. Temperature of solution decreases.
surrounding temperature increases. 5. Container feels cold.
5. Container feels hot. 6. Example:
Example: (a) Ammonium salt dissolves in water.
(a) Neutralisation of acid by alkali
NH NO (s) → NH (aq) + NO (aq)
−
+
4 3 4 3
NaOH(aq) + HCl(aq) → NaCl(aq) + H O(l)
2
(b) Combustion of fuel (b) Hydrated salt is decomposed by heat to form
anhydrous salt.
C H OH(l) + 3O (g) → 2CO (g) + 3H O(l)
2 5 2 2 2
(c) Displacement of copper from copper(II) CuSO .5H O → CuSO + 5H O

4
4
2
2
sulphate solution by zinc metal.
(c) Heat decomposition.
Zn(s) + CuSO (ak) → ZnSO (ak) + Cu(s)
4 4
(d) Precipitation of silver chloride insoluble salt. CaCO → CaO + CO 2

3
AgNO (aq) + HCl(aq) → AgCl(l) + HNO (aq)
3 3
Heat is absorbed
Heat is released


3.2 Heat of Reaction

1. Heat of reaction, ΔH is the change in heat when 1 mole of reactants react or 1 mole of products is formed.
2. When chemical reaction releases heat to the surroundings, ΔH is negative.
3. When chemical reaction absorbs heat from the surroundings, ΔH is positive.
4. Change in energy in a chemical reaction is shown in the energy level diagram.


(a) Energy level diagram for exothermic reactions: (b) Energy level diagram for endothermic reactions:
Energy Energy
Reactants Products
ΔH negative
ΔH positive
Products Reactants



47

PAPER 1


Each question is followed by four options A, B, C or D . Choose the best option for each question.


3.1 Heat Change in Reactions 4. Diagram 1 shows an energy level diagram.
Energy
1. The following thermochemical equation shows Reactants
the combustion of ethanol in oxygen.
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ΔH negative
C H OH + 3O → 2CO + 3H O H = −280 kJ mol −1
2 5 2 2 2
Products
Based on the equation, which statement is
correct? Diagram 1
A The reaction is endothermic. Which of the following is true about the diagram?
−1
B The activation energy is 280 kJ mol . A Heat is absorbed.
C The temperature of the mixture increases. B Endothermic reaction takes place.
D The total energy of the reactants is lower C Temperature of surroundings increases
than the products. during reaction.
D Energy content of the reactants is less than
2. Which of the following is an example of that of the products.
endothermic reaction?
A Solid sodium hydroxide dissolved in distilled 5. Diagram 2 shows the energy profile for a reaction.
water.
Energy
B Solid ammonium nitrate dissolved in
distilled water.
C Dilute hydrochloric acid added to silver R 50 kJ
nitrate solution.
D Dilute hydrochloric acid added to potassium P + Q 200 kJ
hydroxide solution.
Diagram 2
3. Which of the following is correct about What is the activation energy and type of this
exothermic and endothermic reactions? reaction?

Exothermic Endothermic Activation energy (kJ) Type of reaction
reaction reaction A 250 Endothermic

B 250 Exothermic
A Heat is absorbed Heat is released
C 200 Exothermic
B Chemical bonds Chemical bonds D 50 Endothermic
are broken are formed
6. Diagram 3 shows an energy level diagram.
C Temperature of Temperature of
surroundings surroundings Energy
−
increases decreases OH (aq) + H (aq)
+
D Total energy content Total energy content
of the products is of the reactants is H O(l)
2
higher than that higher than that
of the reactants of the products Diagram 3
What conclusion can be made from the diagram?
50 Question 1: Question 5:
SOS TIP Exothermic reaction, ΔH negative Endothermic reaction, ΔH positive

PAPER 2


Section A

Answer all questions.

3
1. Diagram 1 shows the apparatus set-up for an experiment to determine the heat of precipitation. 25 cm
−3
3
−3
of 2.0 mol dm lead(II) nitrate solution is added to 25 cm of 2.0 mol dm sodium sulphate solution in a
polystyrene cups wrapped with hand towels.
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Polystyrene cups wrapped with hand towels

3
−3
+ 25 cm of 2.0 mol dm
3
−3
25 cm of 2.0 mol dm
lead(II) nitrate solution sodium sulphate solution
Diagram 1
(a) What is meant by heat of precipitation from this reaction? [1 mark]



(b) What is the colour of the precipitate formed? [1 mark]




(c) Table 1 shows the results of the experiment. HOTS Analysing

Description Temperature (°C)

Initial temperature of lead(II) nitrate solution 29.0
Initial temperature of sodium sulphate solution 30.0
Highest temperature of the mixture 41.5

Table 1
(i) Mark (3) in the box provided to show which process has the higher heat in the reaction. [1 mark]

Heat absorbed to break the bonds in the reactants.
Heat released during the formation of bonds in the products.

(ii) Calculate the heat energy change in the reaction.
−1
−3
−1
[Specific heat capacity of solution, c = 4.2 J g °C ; density of solution = 1 g cm ] [2 marks]



(iii) Calculate the heat of precipitation for the reaction. [2 marks]







62 Question 1:
SOS TIP (c) Heat energy is absorbed to break existing bonds during reaction.
Heat energy is released when new bonds are formed during reaction.

PAPER 3



1. A student carried out an experiment to determine the heat of combustion of methanol, ethanol, propanol
and butanol. Diagram 1 shows the apparatus set-up for the experiment.
Thermometer

Wind shield

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Copper can Water

Pipeclay triangle


Spirit lamp Fuel
Wooden block
Alcohol
Diagram 1
Table 1.1 shows the mass of lamp before and after burning of the alcohols.

Reading of electronic balance
Alcohol Mass of alcohol used (g)
Before After





Methanol,
CH OH
3
244.95 g 243.40 g






Ethanol,
C H OH
2 5
202.00 g 200.80 g





Propanol,
C H OH
3 7
234.40 g 233.30 g











73
Question 1:
Mass of spirit lamp reduced = Mass of alcohol burnt SOS TIP

4
Chapter Polymer




NOTES



4.1 Polymers (c) After it has set, it cannot alter its shape when
heated and cooled again
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1. A polymer is a long chain molecule that is made from (d) Stronger than thermoplastic
a combination of many repeating basic units known (e) Opaque
as monomer. (f) Can only be moulded once and cannot be
2. Polymerisation is the monomer combination reaction recycled
to produce big, long chained molecules known as Example:
polymers. • Bakelite: Plug, electrical switches, cooking utensil
holder
Covalent bond
Polymerisation • Melamine: Dinnerware, countertops
• Epoxy resin adhesives: Adhesives for metals, glass,
Polymer wood, and battery casing
Monomers 7. Elastomer:
(a) Has few cross linkages.
3. Polymers can be divided into natural polymers and (b) Rubber-like properties, it can be thermoplastic
synthetic polymers.
or thermosetting.
Natural polymers Synthetic polymers (c) More elastic than thermoplastic.
(d) Can be stretched and return to its original shape
Monomer Polymer Monomer Polymer
Example: Natural and synthetic rubber
Glucose Starch Propene Polypropene 8. There are two types of polymerisation: Addition
Isoprene Natural Styrene Polystyrene polymerisation and condensation polymerisation.
rubber 9. During addition polymerisation, alkene monomers
that contain double bonds between carbon atoms,
Amino Protein Ethene Polyethene C=C are added to one another to form polymers
acids
under a high temperature condition.
Example:
4. Based on the intermolecular forces, polymers are
classified into thermoplastic, thermosetting and H H H H
elastomer.
n C C C C
5. Thermoplastic:
(a) Has long chain, no cross linkages. H H H H n
(b) Produced from addition polymerisation. Ethene Polyethene
(c) Can change shape when heated or cooled.
(d) Not as strong as thermosetting. 10. In condensation polymerisation, monomers without
(e) Colourless and transparent. double bonds are combined to form long polymer
(f) Can be moulded repeatedly and can be recycled. chain and by-product.
Example: Example:
• Polyethene: Plastic bag, bottles (a) Formation of nylon-6,10 through condensation
• Polypropylene: Chairs, feeding bottles polymerisation.
• Polystyrene: Food packaging
O O
• Polyvinyl chloride (PVC): Electric wire and cable
H N (CH ) NH + Cl C (CH ) C Cl
insulators, clothes hanger, water pipes 2 2 6 2 2 8
• Perspex: Car windscreen, aircraft windows 1,6-hexanediamine Decanedioyl dichloride
6. Thermosetting:
(a) Has many cross linkages, forms three
dimensional network
(b) Produced from condensation polymerisation
O O
(HN (CH ) (NH C (CH ) C ) + 2HCl
2 6 2 8 n
Nylon-6,10

78

(b) Glucose molecules combined to form starch and releasing simple molecules such as water.
CH OH CH OH CH OH
2 2 2
C O C O C O
HO OH H OH H OH
H H H
C C C C C C
OH H OH H OH H
H H HO H HO H
C C C C C C
H OH H OH H OH
Glucose
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CH OH CH OH CH OH
2 2 2
C O C O C O
H H H H H H H H H
C C C C C C + H O
OH H OH H OH H 2
O O O O
C C C C C C
H OH H OH H OH
Starch
11. Synthetic polymers are very stable. Unlike metals, 7. To produce vulcanised rubber in the school laboratory,
wood or paper, they do not rust, rot or decay easily in strips of rubber are soaked in disulphur dichloride
the presence of water, oxygen or other chemicals or in solution in toluene then dried in air.
sunlight. 8. Vulcanised rubber:
12. Synthetic polymers are widely used in medicine,
C C C C
packaging, coating, textile, electrical and electronic
industry, adhesive, household items and others. S S S S Sulphur
13. Effects of using synthetic polymers on the S S S S cross
C C linkage
environment: C C C C
(a) Synthetic polymers are not biodegradable S S
Rubber
(not decayed by microorganism) so when S S
polymer chain
accumulated, able to block drainage and caused C C
C C C C
flash floods.
(b) When synthetic polymers are burnt, poisonous (a) Sulphur atoms form cross linkages between
gases are produced. the long chains of rubber polymers.
(b) Sulphur cross linkages prevent the long
4.2 Natural Rubber chain of rubber polymers from sliding over
each other easily.
1. Latex is a milky liquid that drips out when incisions (c) Sulphur cross linkages also help the rubber
are made on the bark of the rubber tree trunk, process polymer chains get back to their original
call “tapping”. form after stretching.
2. Latex is a type of coloid that consist of rubber particles (d) Used: to manufacture tyres, rubber hose
and water. and gloves
3. Natural rubber is a natural polymer formed from
monomers called isoprene or 2-methylbut-1,3-diene. Vulcanised rubber Unvulcanised rubber
H CH H H H CH H H Can withstand high Cannot withstand high
3 3
temperatures temperatures
n C C C C C C C C
H H H H n More elastic Less elastic
Isoprene Polyisoprene
(natural rubber) Harder and stronger Less hard, softer
4. Natural rubber polymer is known as polyisoprene has Can resist oxidation Easily oxidised
a formula (C H ) with n valued at as much as 10 000.
5 8 n
5. Latex coagulates in the presence of acid while 4.3 Synthetic Rubber
ammonia solution prevents coagulation of latex.
6. To produce vulcanised rubber in industry, natural 1. Synthetic rubber is man-made polymer produced
rubber is heated with sulphur and zinc oxide as the from polymerisation process of hydrocarbon obtained
catalyst. from fractions of petroleum.


79

2. General properties of synthetic rubber:
(a) Resistant to heat. (d) Hard.
(b) Resistant to chemicals. (e) Heat insulator.
(c) Resistant to oxidation. (f) Elastic.

3. The properties and use of synthetic rubber:
Synthetic rubber Monomer Properties Uses
• Resistant to high heat
Styrene-Butadiene Styrene and • Easily vulcanised Tyres
Rubber (SBR) butadiene, C H Shoe soles
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4 6 • Corrosion resistance
• Elastic
• Resistant to high heat Water pipe, gloves,
Neoprene Chloroprene, C H Cl • Strong tensile strength electrical wire insulators,
4 5
• Resistant to oxidation petrol rubber hose
• Not inflammable
• Less permeable than natural rubber Inner tube of tyres, roofing
• Resistant to heat materials, nonpermeable
Butyl rubber Isobutylene, C H
4 8 • Can be vulcanised with sulphur layer to prevent diffusion
under normal condition of water in reservoir
1,2-dichloroethane, Very resistant against chemical Lining for oils and solvents
Thiokol rubber
C H Cl solvents and oils storage tanks
2 4 2
4. Advantages of synthetic rubber: 5. Disadvantages of synthetic rubber:
(a) Withstand inclement weather and high (a) Ability to absorb vibrations, sound, shock is less
temperatures. compared to natural rubber.
(b) Non-inflammable and can withstand heat. (b) Elasticity and tensile strength are less compared
(c) Good heat and electrical insulators. to natural rubber.
(d) Not easily permeable by gas and water.
(e) Very high resistance towards chemicals.


PAPER 1


Each question is followed by four options A, B, C or D . Choose the best option for each question.

4.1 Polymers



1. Which of the following is the correct pair of 3. Which polymer has the empirical formula of
natural polymer and its monomer? CH ?
2
Natural polymer Monomer
A H H C H H
A Polyisoprene Propene
C C C C
B Starch Amino acid
H C H n H C H n
6 5 2 5
C Cellulose Glucose
D Protein Isoprene B H H D H H
C C C C
2. Which substance is a natural polymer?
H H n H CH 3 n
A Polythene C Polyisoprene
B Polypropene D Polyvinyl chloride

80 Question 2: Question 3:
SOS TIP Polyisoprene is the name of natural rubber. Empirical formula for polymers is the same as its monomer.

PAPER 2


Section A

Answer all questions.

1. The following flow chart shows the change of latex to vulcanised rubber through Steps I and II.

Step I Solid natural Step II Vulcanised
Latex
Acid rubber S Cl 2 rubber
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2
Diagram 1
(a) (i) Step I involves the coagulation of latex to form solid natural rubber. State how Step I is carried out
in the laboratory. [1 mark]




(ii) State two properties of natural rubber. [2 marks]







(b) Step II involves the vulcanisation of rubber. State how Step II is carried out. [3 marks]






(c) (i) Draw the structural formulae of unvulcanised rubber and vulcanised rubber. [2 marks]






















Unvulcanised rubber Vulcanised rubber

(ii) Based on the structural formula in (c)(i), state the structural change that takes place when natural
rubber is vulcanised. [2 marks]






87
Question 1:
(c) (i) Sulphur cross linkages are found in vulcanised rubber.
(c) (ii) Double bonds between carbon atoms in rubber molecule react with sulphur. SOS TIP

PAPER 3



Table 1 shows the result of an experiment to investigate the difference in properties between natural rubber and
vulcanised rubber.



Rubber P Rubber Q

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Initial length (cm) 8.0 cm 8.0 cm







7
Clip Clip
0 8 0
2 2
3 3
9
4 4
5 5
10
Length with 200 g Rubber P 6 Rubber Q 6
7 7
weight (cm) 8 11 8
9 9
10 12 10
11 11
12 12
Weight 200 g Weight 200 g





8.0 cm
Length after weight
is withdrawn (cm) 10.0 cm








Table 1
Based on Table 1, plan an experiment in the laboratory to investigate the elasticity or strength of unvulcanised
and vulcanised rubber.
1. State the variables for this experiment. [2 marks]
(a) Manipulated variable:

(b) Responding variable:





93
• Vulcanisations produce sulphur cross linkages between rubber polymer chains.
• Vulcanisations produce rubber that is more elastic and of better quality. SOS TIP

Answers





0
0
C H AP TE R 1 26. A E value of Zn is more negative than the E value of Fe. Fe is the
CHAPTER 1
cathode and Zn is the anode.
Paper 1 E 0 cell = –0.44 – (–0.76)
= +0.32 V
1. D Acidified potassium dichromate(VI) is an oxidising agent. It 27. A The right-hand side is a chemical cell, S (−) is anode and R (+) is
oxidises alcohols to carboxylic acids. cathode. The left-hand side is an electrolytic cell: P is cathode
2. B Potassium bromide is a reducing agent. Bromide ion is oxidised (−) and Q is anode (+). Reduction occurs at cathodes, R and P.
to bromine. 28. D The right-hand side is a chemical cell: Mg is anode (−); Mg is
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2+
3. A (+3) + 3(x) = 0 oxidised to Mg . The left-hand side is an electrolytic cell: Cu is
2+
3 + 3x = 0 anode (+); Cu is oxidised to Cu .
 x = −1 29. B Cell I: Oxidation of Cu to Cu . Cell II: Oxidation of Cl to Cl .
−
2+
2
4. B Carbon can reduce zinc oxide but cannot reduce aluminium 30. C Anode (P) is plating metal, cathode (Q) is object to be plated
oxide. and electrolyte (R) is silver nitrate.
0
0
5. C KMnO solution acts as an oxidising agent under acidic 31. D E of OH is more negative. OH is oxidised to O gas. E of OH +
–
–
4 2
+
condition. Sulphuric acid supplies the H ions to KMnO . is more positive, H is reduced to H gas.
+
4 2
6. A Combustion of metal in oxygen produces metal oxides are 32. B Higher concentration of Cl , thus Cl ions easier to be oxidised
–
–
redox reactions. at the anode.
7. C A displacement reaction occurred; shows that copper is more 33. B Cells I & III: Hydrogen gas is produced at the cathode and
reactive than metal M. It is more likely to release electrons and chlorine gas is produced at the anode.
be oxidised.
Cell II: Zinc formed at the anode, clorine gas formed at the
8. B Iodide ion is a reducing agent; easily oxidised to iodine. cathode.
9. B Combustion and rusting are redox reactions 34. C Anode: O gas; 4OH → O + 4e + 2H O
−
−
2 2 2
+
−
10. B Cathode: H gas; 2H + 2e → H 2
2
SiO + C ⎯→ Si + CO 2 35. C Hydrogen gas is produced at the cathode and bromine gas is
2
+4 0
+
−
produced at the anode. Na ions combine with OH ions to
11. D Reduction involves loss of oxygen, gain of electrons or gain of
produce sodium hydroxide, NaOH.
hydrogen.
36. B Cell notation of a voltaic cell is written as:
3+
2+
12. B Chlorine water is an oxidising agent that oxidises Fe to Fe . Electrode(s) ǀ Electrolyte(aq) ǁ Electrolyte(aq) ǀ Electrode(s)
Chapter 1
2+
13. C Magnesium displaces Cu ions from copper(II) nitrate. Excess
magnesium powder also observed in the beaker. Electrode Ionic Ionic Electrode
solution solution
14. D 1(+1) + (Cl) + 4(−2) = 0
Cl = +7 Anode (negative terminal) Cathode (positive terminal)
2+
37. B Cu ions are reduced at the cathode to deposit copper metal
15. B A reducing agent releases electrons.
onto the electrode.
2+
2+
16. C Zinc displaces Cu from the solution. When all Cu ions are Cu (aq) + 2e ⎯→ Cu(s)
2+
–
2+
displaced, blue colour disappears and Cu ions change to Cu
38. A Bauxite is an aluminium ore. Aluminium is extracted by
atoms, producing a brown solid.
electrolysis of molten bauxite.
17. A Al and Ca are above Ag in the reactivity series of metals.
39. A Metals above carbon in the reactivity series of metals are
Sulphur is a non-metal.
extracted by electrolysis of their molten ores. Carbon is
−
18. B Chlorine, Cl is an oxidising agent; oxidises I to I . Iodine
2 2 positioned between aluminium and zinc.
solution is brown.
40. C Carbon reduces oxides of zinc and copper. Zinc reduces oxides
19. D Acidified KMnO is an oxidising agent. SO and H S
4 2 2 of iron and copper. Electrolysis is used in purification of metal.
are reducing agents that will react with KMnO , CH is
4 4 41. C Ca is above Mg in the reactivity series of metals.
hydrocarbon.
42. D Iron(III) oxide is reduced by carbon and carbon monoxide to
20. B The chemical formula of chromium(VI) oxide is CrO .
3
iron and carbon dioxide.
21. D A more reactive metal reduces the oxide of a less reactive
43. D Al is formed at the cathode by reduction. Al ions gain
3+
metal. Q > R > S > P
electrons to form Al atoms.
22. B Graphite and platinum are inert electrodes.
44. C Carbon (coke) is more reactive than iron but less reactive
23. D The more positive the E° value, the easier reduction occurs at
than aluminium. Carbon (coke) reduces iron oxide but not
the cathode (positive terminal). aluminium oxide.
E 0 = (+0.34) – (−0.44) = + 0.78 V
cell
45. A Cryolite has a low melting point than bauxite. It reduces the
24. B The more negative the E° value, the more readily for oxidation
high melting point of bauxite thus save energy.
to occur; stronger reduction agent.
46. D A sacrificial metal must be more electropositive than iron.
0
25. D The more positive the E value, the more readily for reduction
A sacrificial metal lose its electrons easier. Sacrificial metal
to occur; stronger oxidation agent.
ionises before iron, corrosion of iron is prevented.
114

2+
2+
47. B Copper is the least electropositive metal that has the slowest 4. (a) Zn(s) + Cu (aq) → Zn (aq) + Cu(s)
rate of corrosion. (b) (i) and (ii)
48. B Magnesium is more electropositive than iron. Manesium Voltmeter
release electrons more easily than iron. Magnesium undergoes
oxidation. Rusting of iron is prevented.
Anode Cathode
49. A Aluminium has a protective oxide layer against corrosion.
Aluminium and aluminium oxide are amphoteric. Electrode 2 Copper
electrode
50. D Sodium is a reactive metal and reacts vigorously with air and Copper(II)
sulphate
water. solution

(c)
Voltmeter reading (V) Metal
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Paper 2
1.10 Zinc
Section A 0.78 Iron
1. (a) Electrolysis is a process to decompose a compound into its 0.38 Tin
elements by passing an electric current through the compound 0.00 Copper
in its molten or aqueous state.
(d) Magnesium
(b) Acid X: Hydrochloric acid, HCl(aq)
5. (a) Electrolysis of water
Acid Y: Nitric acid, HNO (aq)
3
(b) Sodium hydroxide
(c) Chorine gas
(c) (i) and (ii)
−
(d) (i) 2H (aq) + 2e → H (g)
+
2 Anode: Oxidation (loss of electrons)
(ii) 2H O(l) → O (g) + 4H (aq) + 4e −
+
2 2 Cathode: Reduction (gain of electrons)
(e) (i) Use a lower concentration of hydrochloric acid. The higher (d) The product is water, a non-toxic substance
the concentration of Cl ions, the easier for Cl ion to be
−
6. (a) Haematite
oxidised to chlorine gas.
(b) (i) Limestone, CaCO
−
2
(ii) 2Cl (aq) → Cl (g) + 2e − 3
(ii) To remove calcium silicate (slag) through reaction
2. (a) Electrolyte is a compound that conducts an electric current in
between silicon dioxide and calcium oxide.
the molten or aqueous state and undergo chemical changes.
CaO + SiO → CaSiO (slag)
(b) (i) There is a flow of current/electrons in the circuit. 2 3 Chapter 1
(c) (i) Carbon monoxide
2+
−
(ii) Copper(II) ion, Cu and chloride ion, Cl .
(ii) C + CO → 2CO
(iii) Yellow-green gas is liberated. 2
7. (a) (i) Fe 2+
(iv) 2Cl (l) → Cl (g) + 2e −
−
2 (ii) Fe 3+
(c) (i) The electrode connected to the negative terminal of the
(b) (i) P: Magnesium/Aluminium/Zinc [any one]
battery.
2+
(ii) Mg(s) + Fe (aq) → Mg (aq) + Fe(s)
2+
(ii) Brown solid deposited on the cathode.
(c) (i) Q: Carbon
(iii) Cu (aq) + 2e → Cu(s)
−
2+
(ii) 3C(s) + 2Fe O (s) → 3CO (g) + 4Fe(s)
3. (a) (i) A redox reaction involves oxidation and reduction taking 2 3 2
place simultaneously.
(ii) C(s) + 2PbO(s) → 2Pb(s) + CO (g) Section B
2
(iii) +2 to 0 8. (a) Oxidising agent: Hydrogen peroxide
(iv) Lead(II) oxide Reducing agent: Iodide ion
(b) Carbon is more reactive than metal X. (b) Oxidation: 2I (aq) → I (aq) + 2e
−
−
2
Carbon is less reactive than zinc. Reduction: H O (aq) + 2H (aq) + 2e → 2H O(l)
+
−
2 2 2
(c) (i) Zinc, carbon, lead, X (c) Iodine: −1 to 0
(ii) X is copper Oxygen: −1 to -2
Hydrogen: no change, +1
(d) Crucible
(d) Colourless solution turns brown
Metal oxide +
carbon (e) Product of reduction of hydrogen peroxide is water.
Water is a non-toxic substance.
9. (a) Solution R: Iron(II) sulphate solution
Solution S: Potassium iodide solution
(b) (i) Set I: Oxidising agent: Acidified potassium
dichromate(VI)
Reducing agent: Iron(II) sulphate solution
Set II: Oxidising agent: Chlorine water
Reducing agent: Potassium iodide solution
115

2+
(ii) Set I: Iron(II) ion, Fe Discussion:
Set II: Iodide ion, I − Potassium hexacyanoferrate(III) solution confirms the
(iii) Set I: Fe (aq) → Fe (aq) + e presence of Fe ions which are produced when iron rusts.
2+
−
2+
3+
−
Set II: 2I (aq) → I (aq) + 2e − The more the dark blue spots formed, the higher the
2
corrosion of iron. Phenolphthalein detects the presence of
OH ions which are produced at the cathode during the
−
Section C corrosion of metals. The higher the intensity of the pink
colour, the greater the extent of rusting of iron. A metal more
10. (a) Iron undergoes oxidation at the anode:

2+
Fe(s)→ Fe (aq) + 2e − electropositive than iron can protect the iron from rusting. A
less electropositive metal than iron speeds up the corrosion
Electrons flow through iron to the cathode (edge of water
of iron.
droplet). Oxygen is reduced to hydroxide ion, OH .
−
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−
O + 2H O + 4e → 4OH
2 2 11. (a) (i) Solution X in Set I: Dilute hydrochloric acid
−
Fe ions combined with OH ions to form a green precipitate, At the anode, the E value of OH is less positive that the
2+
0
–
iron(II) hydroxide. E value of Cl . Oxidation of water to O gas occurs more
0
–
2
Fe (aq) + 2OH (aq) → Fe(OH) (s) readily.
−
2+
2
Further oxidation by oxygen, produced brown precipitate, Oxygen gas is colourless.
hydrated iron(III) oxide or rust. Solution X in Set II: Concentrated hydrochloric acid
oxidation
Fe(OH )(s) ⎯⎯⎯→ Fe O .xH O(s) At the anode, the concentration of Cl ion is higher than
−
2 2 3 2
–
−
(b) (i) Metal X: Magnesium OH ion. Oxidation of Cl to Cl gas occurs more readily.
2
Metal Y: Lead Chlorine gas is yellow-green.
−
+
Magnesium is more electropositive than iron. (ii) H ions and Cl ions
Magnesium is more likely to release electrons compared (iii) Set I: Anode: 4OH (aq) → O (g) + 2H2O(l) + 4e −
−
2
to iron and oxidised. Cathode: 2H (aq) + 2e → H (g)
−
+
2+
Mg(s) → Mg (aq) + 2e − 2
Set II: Anode: 2Cl (aq) → Cl (g) + 2e −
−
2
Magnesium preventing iron from rusting. While lead is Cathode: 2H (aq) + 2e → H (g)
−
+
2
less electropositive than iron. Iron releases electrons and
(b) Metal P: Iron (Fe); Metal Q: Copper (Cu); Metal R: Zinc (Zn)
oxidised. Iron will rusts. P: Iron(II) nitrate; Q: Copper(II) nitrate; R: Zinc nitrate
2+
Fe(s) → Fe (aq) + 2e −
(i) Materials: Zinc foil, iron foil, iron(II) nitrate solution,
(ii) Procedure:
copper(II) nitrate solution
Test tube I Test tube II Test tube III Apparatus: Beaker
(ii) Procedure:
Iron nail Iron nail Iron nail
I II III
Chapter 1
coiled with coiled with coiled with
zinc strip copper strip aluminium strip Zinc foil Iron foil
Agar solution + potassium hexacyanoferrate(III)
Beaker
+ phenolphthalein indicator
Test tube IV Test tube V Iron(II) nitrate Copper(II) nitrate
solution solution
Iron nail Iron nail 1. Set up the apparatus as shown above with a metal dipped
coiled with into a nitrate salt solution.
tin strip
2. Observe the changes that occur in each beaker.
1. Iron nails together with metal strips are cleaned using
Results:
sandpaper.
Beaker I: Zinc dissolves; grey solid deposited; green
2. Four iron nails are each coiled with different metals.
solution turns colourless.
3. The five iron nails are placed into separate test tubes
containing agar solution mixed with potassium Beaker II: Zinc dissolves; brown solid deposited; blue
hexacyanoferrate(III) and phenolphtalein. solution turns colourless.
4. The test tubes are left in a test tube rack for two days. Beaker III: Iron dissolves; brown solid deposited; blue
5. All changes are recorded. solution turns colourless.
Observation: Conclusion: Zinc is the most electropositive metal. It displaces
iron and copper from their nitrate salt solutions.
Test Dark Intensity of Inference
tube blue spot pink colour Iron is more electropositive than copper. It
displaces copper from its salt solution.
I None High Iron nail did not rust
Cu, Fe, Zn
II Most Low Iron nail rusted the most
more reactive
III None High Iron nail did not rust
IV Moderate Low Iron nail rusted a bit
V Few - Iron nail rusted a bit
116

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