Contents Must Know iii – xii Chapter 1 Redox Equilibrium 1 – 21 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
Mnemonics (Chapter 1) 1 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 1) 3 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 1) 5 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 1) 7 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 2) 9 @ Pan Asia Publications Sdn. Bhd. Mnemonics (Chapter 3) 11 @ Pan Asia Publications Sdn. Bhd. Redox Equilibrium O Oxidation I Is L Loss R Reduction I Is G Gain Thermochemistry Voltaic Cell and Electrolytic Cell FAT CAT Electrons flow From Anode To CAThode. Electrolytic Cell • Cation = Ca+ion 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) • OxidAtion occurs at Anode • ReduCtion occurs at Cathode Electrolyte e– e– V Voltaic cell EXothermic reaction Heat EXit from system ENdothermic reaction Heat ENter the system Heat Heat To Determine Oxidising Agents and Reducing Agents Based on the Value of E0 Molecules or ions with a more positive (less negative) E0 value: AcRO Ac Accept electron R Reduction reaction O Oxidising agent Atoms or ions with a more negative (less positive) E0 value: DoOR Do Donate electron O Oxidation reaction R Reducing agent Prefix for Naming Carbon Compounds Meth Monkey Eth Eats Prop Pile But Bananas Monkey Eats a Pile of Bananas Battery e– e– + – + – Anode Electrolyte Cathode – Anion + Cation Electrolytic cell KNOW Mnemonics
Important Definitions (Chapter 4) 20 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 5) 22 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 5) 24 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 3) 14 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 3) 16 @ Pan Asia Publications Sdn. Bhd. Important Definitions (Chapter 3) 18 @ Pan Asia Publications Sdn. Bhd. Fats or oils Fatty acid + Glycerol Fatty acid salt (soap) Precipitation Thermoplastic, Thermosetting and Elastomer • Thermoplastics are loose polymer that can be heated and reshaped repeatedly. • Thermosettings are rigid polymer that will not return to their original shape when heated. • Elastomers are polymer that can be deformed and returned to their original shape. Soap and Detergent • Soaps are sodium or potassium salts of fatty acids. • Detergents are sodium salts of sulphonic acids or alkyl hydrogen sulphates. • Saponification is the process to prepare soaps by alkaline hydrolysis of fats or oils. • Nanoscience is a study on processing of substances at nanoscale, between 1 nanometre - 100 nanometres. • Nanotechnology is the study and manipulation of matter at nanometer scale to produce new materials or devices. • Graphene is the carbon allotrope. Graphene is one-layer thick graphite arranged in hexagonal honeycomb-like structure to form graphene sheets. • Graphene sheets can used to produce graphites, carbon nanotubes and fullerene balls. Types of Chemical Reaction Exothermic reaction: • Chemical reaction that releases heat energy to the surroundings. • Temperature of the surroundings increases. Endothermic reaction: • Chemical reaction that absorbed heat energy from the surroundings. • Temperature of the surroundings decreases. Heat of Displacement • Heat of displacement is the heat change when one mole of a metal is displaced from its aqueous salt solution by a more electropositive metal. • Heat of displacement = − Q n Heat change, Q = mcθ = V × 4.2 J g−1 °C−1 × θ Number of moles, n = MV 1 000 Heat of Combustion • Heat of combustion is the total heat released when one mole of fuel is burnt completely in excess oxygen gas. • Heat of combustion is given in the unit kJ mol−1. • Heat of combustion = − Q n Heat change, Q = mcθ = V × 4.2 J g−1 °C−1 × θ Number of moles, n = Mass Molar mass KNOW Important Definitions
Important Diagrams (Chapter 1) 49 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 2) 51 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 3) 53 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 4) 55 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 5) 57 @ Pan Asia Publications Sdn. Bhd. Important Diagrams (Chapter 5) 59 @ Pan Asia Publications Sdn. Bhd. Battery e– e– + – + – Anode Electrolyte Cathode – Anion + Cation Cathode • Attracts cations • Negative electrode • Reduction occurs Anode • Attracts anions • Positive electrode • Oxidation occurs Energy Reactants Products ΔH positive Ethene gas Water Combustion tube Glass wool soaked in propanol Porcelain chips Glass wool Heat soaked in ethanol Heat the porcelain chips first before heating the glass wool soaked in ethanol gently. Energy Level Diagram of Endothermic Reaction Apparatus Set up for Dehydration of Ethanol Electrolytic Cell Formation of Grease Droplets During the Cleansing Action of Soap The Vulcanisation Process of Natural Rubber C S C C C C C C C C C C C C C C C C C C C C C C S C C S C C S C S C S C C C C C C C S C S C Vulcanisation Vulcanised rubber Natural rubber Preparation of Detergents Alkylbenzene Alkyl benzene sulphonic acid Concentrated H2 SO4 NaOH Alkyl benzene sulphonic acid + water Sodium alkyl benzene sulphonate + water Step 1 Step 2 Sulphonation Neutralisation + + Soap anions Hydrophobic part Hydrophilic part Droplet of grease KNOW Important Diagrams
1 Chapter 1 Redox Equilibrium 1.1 Oxidation and Reduction 1. Definition of oxidation and reduction. Oxidation Reduction • Gain of oxygen • Loss of oxygen • Loss of hydrogen • Gain of hydrogen • Loss of electron • Gain of electron • Increase in oxidation number • Decrease in oxidation number 2. Oxidation number or oxidation state is the charge of the element in a compound if the transfer of electrons occurs in an atom to form chemical bonds with other atoms. 3. Rules for assigning oxidation number: (a) The oxidation number of each atom of a free element is 0. (b) The oxidation number of a monoatomic ion is equal to the charge of the ion. (c) Fluorine in its compounds has a fixed oxidation number of −1. (d) Alkali metals (Group 1 elements) in their compounds have a fixed oxidation number of +1. (e) Alkaline earth metals (Group 2 elements) in their compounds have a fixed oxidation number of +2. (f) Hydrogen in a compound normally has an oxidation number of +1 when it combines with non-metals, but hydrogen has an oxidation number of −1 when it combines with metals hydrides. (g) Oxygen in a compound normally has an oxidation number of −2 except peroxide compounds (oxidation number is −1) and compounds with fluorine (oxidation number is +2). (h) Halogens (Cl, Br, I) in their compounds have oxidation number of −1 except when combined with fluorine and oxygen. (i) The sum of oxidation numbers of all atoms in a neutral compound is always 0. 4. A redox reaction is a chemical reaction in which oxidation and reduction occur simultaneously. 5. An oxidising agent is a substance that oxidises another substance and reduces itself. Example: (a) Acidified potassium manganate(VII) solution: MnO4 −(aq) + 8H+(aq) + 5e− → Mn2+(aq) + 4H2 O(l) (b) Acidified potassium dichromate(VI) solution: Cr2 O7 2−(aq) + 14H+(aq) + 6e− → 2Cr3+(aq) + 7H2 O(l) (c) Chlorine water: Cl2 (aq) + 2e− → 2Cl−(aq) (d) Bromine water: Br2 (aq) + 2e− → 2Br−(aq) 6. A reducing agent is a substance that reduces another substance and oxidises itself. Example: (a) Magnesium: Mg(s) → Mg2+(aq) + 2e- (b) Sulphur dioxide: SO2 (aq) + 2H2 O(l) → SO4 2−(aq) + 4H+(aq) + 2e− (c) Iron(II) ion: Fe2+(aq) → Fe3+(aq) + e− 7. Examples of redox reactions: (a) Conversion of Fe2+ ion to Fe3+ ion Ionic equation: Br2 (aq) + 2Fe2+(aq) → 2Br−(aq) + 2Fe3+(aq) Reduction half-equation: Br2 (aq) + 2e− → 2Br−(aq) Oxidation half-equation: Fe2+(aq) → Fe3+(aq) + e− (b) Displacement of copper by zinc from copper(II) sulphate solution Ionic equation: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) Reduction half-equation: Cu2+(aq) + 2e− → Cu(s) Oxidation half-equation: Zn(s) → Zn2+(aq) + 2e− (c) Displacement of bromine by chlorine from potassium bromide solution Ionic equation: Cl2 (aq) + 2Br−(aq) → 2Cl−(aq) + Br2 (aq) Reduction half-equation: Cl2 (aq) + 2e− → 2Cl−(aq) Oxidation half-equation: 2Br−(aq) → Br2 (aq) + 2e− 8. Displacement of metal from its salt solution (a) A more electropositive metal displaces a less electropositive metal from its salt solution. (b) The electropositivity of a metal is determined from its position in the electrochemical series. (c) A metal can displace any metal below it from its salt solution. NOTES
2 (d) Three observations may be observed during a displacement reaction of a metal. • The more reactive metal dissolves • The less reactive metal is deposited • The colour of the salt solution may change 9. Displacement of halogen from its halide solution (a) Halogens are oxidising agents. Conversely, halide ions are reducing agents. (b) A halogen molecule, X2 , gains electrons to form the halide ion, X− : X2 (aq) + 2e− → 2X−(aq) (c) Reactivity or oxidising power of the halogens decreases going down Group 17. (d) The halide ion gains electrons to form the halogen molecule, X2 : 2X−(aq) + 2e− → X2 (aq) (e) Reducing power of the halide ions increases moving down Group 17. 1.2 Standard Electrode Potential 1. Standard electrode potential, E0 is defined as the difference of electrode potential (voltage) of an electrode system consisting of an electrode half-cell pairing up with the standard hydrogen electrode (SHE) half-cell. Voltmeter Salt bridge Platinum electrode 298 K and 1 atm H2 (g) Electrode X Solution of metal X ion 1.0 mol dm–3 Acid solution (concentration of H+ is 1.0 mol dm–3 ) V SHE 2. The standard condition to measure the standard electrode potential, E0 of the cell: (a) Concentration of ion in aqueous solution is 1.0 mol dm−3. (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. 3. The E0 value of standard hydrogen electrode, SHE = 0.00 V. H+(aq) + e– ⇌ 1 2 H2 (g) E0 = 0.00 V 4. The above cell can be represented in the form of a cell notation. Pt(s) | H2 (g) | H+(aq) || Xn+(aq) | X(s) ⎧ ⎪ ⎪ ⎨ ⎪ ⎪ ⎩ ⎧ ⎪ ⎨ ⎪ ⎩ • Represents the SHE • Represents electrode X 5. The standard cell potential can be obtained based on the following equation. E0 cell = E0 cathode − E0 anode 6. The electrode potential value, E0 is used to predict: (a) Substance that undergoes oxidation or reduction. (b) Substance that acts as an oxidising or reducing agent. (c) Strength of oxidising or reducing agents. Example: F2 + 2e− 2F− E0 = +2.87 V CI2 + 2e− 2CI− E0 = +1.36 V Br2 + 2e− 2Br− E0 = +1.07 V • E° value increases, strength as oxidising agent increases. • Increasing order of oxidising agent strength: Br2 , Cl2 , F2 Oxidising agent 1.3 Voltaic Cell 1. A simple voltaic cell consists of two different metals immersed in an electrolyte solution and connected with connecting wires. 2. A Daniell cell is an example of a voltaic cell where zinc metal and copper metal are used as electrodes dipped into their respective ionic salt solutions. (+) Voltmeter Zn ZnSO4 CuSO4 Cu KCl e– e– (–) Cathode (Reduction) Anode (Oxidation) Salt bridge Zn2+ + 2e− Zn E0 = −0.76 V Cu2+ + 2e− Cu E0 = +0.34 V Zn(s) → Zn2+(aq) + 2e− (Oxidation) Cu2+(aq) + 2e− → Cu(s) (Reduction) Overall ionic equation: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) Cell notation of Daniell cell: Zn(s) | Zn2+(aq, 1.0 mol dm−3) || Cu2+(aq, 1.0 mol dm−3) | Cu(s) Cell voltage, E0 cell = E0 cathode − E0 anode = E0 Cu − E0 Zn = 0.34 − (−0.76) = +1.10 V
4 SOS TIP Anode (–) Fe(s) Fe2+(aq) + 2e− Iron metal O2 OH− O2 Drop of water Rust (Fe2 O3 .xH2 O) Fe2+ e− Cathode (+) O2 (g) + 2H2 O(l) + 4e− 4OH−(aq) 2. Corroded iron is covered with a layer of red-brown compound. The red-brown compound is hydrated iron(III) oxide, Fe2 O3 .xH2 O or rust. 3. Corroded copper is covered with a green layer. The green compound is basic copper(II) carbonate, CuCO3 . 4. Ways to prevent rusting: (a) Use a protective surface • A protective surface prevents water and air from coming into contact with the iron. Example: paint, grease, oil, plastic, metals and using other metals. • Example of using other metals: Galvanisation involves coating iron with a layer of zinc. Tin plating or tinning is a process of coating iron with a thin layer of tin. Chrome plating is used on car bumpers and decorative items. (b) Alloying • Alloy such as stainless steel is made of a mixture of iron with chromium and nickel. This alloy has a very high resistance to corrosion. Chromium and nickel form oxide layer that is very strong and impermeable to air and water. (c) Sacrificial protection • Involves placing a more electropositive metal in contact with iron. • Example: Magnesium acts as a sacrificial metal to protect underground iron pipes from rusting. Each question is followed by four options A, B, C or D . Choose the best option for each question. PAPER 1 1.1 Oxidation and Reduction 1. Ethanol is the second member of the alcohol family. Ethanol reacts with acidified potassium dichromate(VI) to form ethanoic acid. What type of reaction is involved in this conversion? A Esterification C Halogenation B Dehydration D Oxidation 2. Oxidising agent is used to oxidise iron(II) sulphate to iron(III) sulphate. Which of the following is not an oxidising agent for the reaction? A Chlorine B Potassium bromide C Hydrogen peroxide D Concentrated nitric acid 3. What is the oxidation number of hydrogen in aluminium hydride, AlH3 ? A −1 C +1 B 0 D +2 CLONE SPM CLONE SPM CLONE SPM 4. Carbon is a non-metal and it can be used to reduce metal oxides. In the reactivity series of metals, carbon is located between A calcium and aluminium B aluminium and zinc C zinc and iron D iron and lead 5. Diagram 1 shows the experimental set-up used to investigate a redox reaction. During a redox reaction, transfer of electron at a distance takes place between the oxidising agent and reducing agent which were separated by solution X. Graphite electrodes Potassium iodide solution Potassium manganate (VII) solution Solution X G – + Diagram 1 CLONE SPM CLONE SPM Question 4: Refer to the positions of metals in the reactivity series of metals. You need to memorise this series. Question 5: Reduction half-equation: MnO4 − + 8H+ + 5e− → Mn2+ + 4H2 O H+ ions are supplied by aqueous acid solution HOTS Analysing
10SOS TIP 47. Diagram 16 shows an experiment to investigate the rate of corrosion of different metals. I II Distilled water Iron Copper III IV Distilled water Magnesium Lead Diagram 16 In which test tube will the rate of corrosion be the slowest? A I B II C III D IV 48. Underground iron pipes are connected to a block of magnesium as shown in Diagram 17. Magnesium block Underground iron pipe Diagram 17 Why does the magnesium block stops the iron from rusting? A Magnesium forms an alloy with iron B Magnesium reacts more readily than iron C Magnesium prevents oxygen from coming into contact with iron D Magnesium forms a protective coating of magnesium oxide on the iron surface 49. Diagram 18 shows the reactions of metal P. HOTS Analysing Colourless solution Metal P P oxide Corrodes slowly Acid Alkali Diagram 18 What is metal P? A Aluminium C Calcium B Iron D Copper 50. A method to protect iron from rusting is to apply a protective coating as shown in Diagram 19. HOTS Analysing Coating of Iron substance X Diagram 19 Which of the following could not be X? A Aluminium C Magnesium B Gold D Sodium 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. Graphite electrodes Graphite electrode Acid X Acid Y W X Y Z Graphite electrode Cell I Cell II Diagram 1 PAPER 2
20SOS TIP PAPER 3 The standard electrode potential, E0 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. Salt bridge H2 (g) 1 atm; 25 °C Voltmeter Electrode Acid solution Platinum electrode 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 system Anode Cathode Volmeter diagram Volmeter reading (V) SHE + Mg Mg SHE 0 1 2 5 4 3 SHE + Ag SHE Ag 0 1 2 5 4 3 SHE + Cu SHE Cu 0 1 2 5 4 3 SHE + V V SHE 0 1 2 5 4 3 SHE + Ti Ti SHE 0 1 2 5 4 3 Table 1 Question 1: The E0 value of standard hydrogen electrode, SHE is 0.00 V.
47 Chapter 3 Thermochemistry 3.1 Heat Change in Reactions Exothermic reactions Endothermic reactions 1. Reaction that releases heat energy to the surroundings. 2. Energy content of reactants is higher than that of products. 3. ΔH has a negative value. 4. Heat released to the surroundings, the surrounding temperature increases. 5. Container feels hot. Example: (a) Neutralisation of acid by alkali NaOH(aq) + HCl(aq) → NaCl(aq) + H2 O(l) (b) Combustion of fuel C2 H5OH(l) + 3O2 (g) → 2CO2 (g) + 3H2 O(l) (c) Displacement of copper from copper(II) sulphate solution by zinc metal. Zn(s) + CuSO4(ak) → ZnSO4(ak) + Cu(s) (d) Precipitation of silver chloride insoluble salt. AgNO3(aq) + HCl(aq) → AgCl(l) + HNO3(aq) 1. Reaction that absorbs heat energy from the surroundings. 2. Energy content of products is higher than that of reactants. 3. ΔH has a positive value. 4. Temperature of solution decreases. 5. Container feels cold. 6. Example: (a) Ammonium salt dissolves in water. NH4 NO3(s) → NH4 +(aq) + NO3 −(aq) (b) Hydrated salt is decomposed by heat to form anhydrous salt. CuSO4 .5H2 O → CuSO4 + 5H2 O (c) Heat decomposition. CaCO3 → CaO + CO2 Heat is released Heat is absorbed 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: Energy Reactants Products ΔH negative (b) Energy level diagram for endothermic reactions: Energy Reactants Products ΔH positive NOTES
50SOS TIP Each question is followed by four options A, B, C or D . Choose the best option for each question. PAPER 1 3.1 Heat Change in Reactions 1. The following thermochemical equation shows the combustion of ethanol in oxygen. C2 H5OH + 3O2 → 2CO2 + 3H2 O H = −280 kJ mol−1 Based on the equation, which statement is correct? A The reaction is endothermic. B The activation energy is 280 kJ mol−1. C The temperature of the mixture increases. D The total energy of the reactants is lower than the products. 2. Which of the following is an example of endothermic reaction? A Solid sodium hydroxide dissolved in distilled water. B Solid ammonium nitrate dissolved in distilled water. C Dilute hydrochloric acid added to silver nitrate solution. D Dilute hydrochloric acid added to potassium hydroxide solution. 3. Which of the following is correct about exothermic and endothermic reactions? Exothermic reaction Endothermic reaction A Heat is absorbed Heat is released B Chemical bonds are broken Chemical bonds are formed C Temperature of surroundings increases Temperature of surroundings decreases D Total energy content of the products is higher than that of the reactants Total energy content of the reactants is higher than that of the products 4. Diagram 1 shows an energy level diagram. Energy Reactants Products ΔH negative Diagram 1 Which of the following is true about the diagram? A Heat is absorbed. B Endothermic reaction takes place. C Temperature of surroundings increases during reaction. D Energy content of the reactants is less than that of the products. 5. Diagram 2 shows the energy profile for a reaction. P + Q R 50 kJ 200 kJ Energy Diagram 2 What is the activation energy and type of this reaction? Activation energy (kJ) Type of reaction A 250 Endothermic B 250 Exothermic C 200 Exothermic D 50 Endothermic 6. Diagram 3 shows an energy level diagram. Energy OH−(aq) + H+(aq) H2 O(l) Diagram 3 What conclusion can be made from the diagram? Question 1: Exothermic reaction, ΔH negative Question 5: Endothermic reaction, ΔH positive
62SOS TIP Section A Answer all questions. 1. Diagram 1 shows the apparatus set-up for an experiment to determine the heat of precipitation. 25 cm3 of 2.0 mol dm−3 lead(II) nitrate solution is added to 25 cm3 of 2.0 mol dm−3 sodium sulphate solution in a polystyrene cups wrapped with hand towels. Polystyrene cups wrapped with hand towels + 25 cm3 of 2.0 mol dm−3 lead(II) nitrate solution 25 cm3 of 2.0 mol dm−3 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. [Specific heat capacity of solution, c = 4.2 J g−1 °C−1; density of solution = 1 g cm−3] [2 marks] (iii) Calculate the heat of precipitation for the reaction. [2 marks] PAPER 2 Question 1: (c) Heat energy is absorbed to break existing bonds during reaction. Heat energy is released when new bonds are formed during reaction.
73SOS TIP 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. Wind shield Fuel Wooden block Spirit lamp Copper can Water Thermometer Pipeclay triangle Diagram 1 Table 1.1 shows the mass of lamp before and after burning of the alcohols. Alcohol Reading of electronic balance Mass of alcohol used (g) Before After Methanol, CH3 OH 244.95 g 243.40 g Ethanol, C2 H5 OH 202.00 g 200.80 g Propanol, C3 H7 OH 234.40 g 233.30 g Question 1: Mass of spirit lamp reduced = Mass of alcohol burnt Alcohol
78 Chapter 4 Polymer 4.1 Polymers 1. A polymer is a long chain molecule that is made from a combination of many repeating basic units known as monomer. 2. Polymerisation is the monomer combination reaction to produce big, long chained molecules known as polymers. Monomers Polymer Covalent bond Polymerisation 3. Polymers can be divided into natural polymers and synthetic polymers. Natural polymers Synthetic polymers Monomer Polymer Monomer Polymer Glucose Starch Propene Polypropene Isoprene Natural rubber Styrene Polystyrene Amino acids Protein Ethene Polyethene 4. Based on the intermolecular forces, polymers are classified into thermoplastic, thermosetting and elastomer. 5. Thermoplastic: (a) Has long chain, no cross linkages. (b) Produced from addition polymerisation. (c) Can change shape when heated or cooled. (d) Not as strong as thermosetting. (e) Colourless and transparent. (f) Can be moulded repeatedly and can be recycled. Example: • Polyethene: Plastic bag, bottles • Polypropylene: Chairs, feeding bottles • Polystyrene: Food packaging • Polyvinyl chloride (PVC): Electric wire and cable insulators, clothes hanger, water pipes • Perspex: Car windscreen, aircraft windows 6. Thermosetting: (a) Has many cross linkages, forms three dimensional network (b) Produced from condensation polymerisation NOTES (c) After it has set, it cannot alter its shape when heated and cooled again (d) Stronger than thermoplastic (e) Opaque (f) Can only be moulded once and cannot be recycled Example: • Bakelite: Plug, electrical switches, cooking utensil holder • Melamine: Dinnerware, countertops • Epoxy resin adhesives: Adhesives for metals, glass, wood, and battery casing 7. Elastomer: (a) Has few cross linkages. (b) Rubber-like properties, it can be thermoplastic or thermosetting. (c) More elastic than thermoplastic. (d) Can be stretched and return to its original shape Example: Natural and synthetic rubber 8. There are two types of polymerisation: Addition polymerisation and condensation polymerisation. 9. During addition polymerisation, alkene monomers that contain double bonds between carbon atoms, C=C are added to one another to form polymers under a high temperature condition. Example: H H C H H C H Ethene Polyethene H n C H n H C 10. In condensation polymerisation, monomers without double bonds are combined to form long polymer chain and by-product. Example: (a) Formation of nylon-6,10 through condensation polymerisation. (CH2 (CH2 )8 )6 1,6-hexanediamine Decanedioyl dichloride Nylon-6,10 H2 N NH2 +Cl C Cl O C O (CH + 2HCl 2 (CH2 )8 )6 (HN (NH ) C n O C O
79 (b) Glucose molecules combined to form starch and releasing simple molecules such as water. O C H CH2 OH C H OH C OH H H HO C H OH C O C H CH2 OH C H OH C OH H HO H C H OH C O C H CH2 OH C H OH C OH H HO H C H OH C Glucose O C H CH2 OH C H OH C OH H O H O H H C C O C H CH2 OH C H OH C OH H C C O C H CH2 OH C H OH C OH H C C O H H O H + H2 O Starch 11. Synthetic polymers are very stable. Unlike metals, wood or paper, they do not rust, rot or decay easily in the presence of water, oxygen or other chemicals or in sunlight. 12. Synthetic polymers are widely used in medicine, packaging, coating, textile, electrical and electronic industry, adhesive, household items and others. 13. Effects of using synthetic polymers on the environment: (a) Synthetic polymers are not biodegradable (not decayed by microorganism) so when accumulated, able to block drainage and caused flash floods. (b) When synthetic polymers are burnt, poisonous gases are produced. 4.2 Natural Rubber 1. Latex is a milky liquid that drips out when incisions are made on the bark of the rubber tree trunk, process call “tapping”. 2. Latex is a type of coloid that consist of rubber particles and water. 3. Natural rubber is a natural polymer formed from monomers called isoprene or 2-methylbut-1,3-diene. H H C H C Isoprene Polyisoprene (natural rubber) n C H n H C H H C H C H H C CH3 C CH3 4. Natural rubber polymer is known as polyisoprene has a formula (C5 H8 )n with n valued at as much as 10 000. 5. Latex coagulates in the presence of acid while ammonia solution prevents coagulation of latex. 6. To produce vulcanised rubber in industry, natural rubber is heated with sulphur and zinc oxide as the catalyst. 7. To produce vulcanised rubber in the school laboratory, strips of rubber are soaked in disulphur dichloride solution in toluene then dried in air. 8. Vulcanised rubber: Sulphur cross linkage Rubber polymer chain C S S C C S S C C C C S S C (a) Sulphur atoms form cross linkages between the long chains of rubber polymers. (b) Sulphur cross linkages prevent the long chain of rubber polymers from sliding over each other easily. (c) Sulphur cross linkages also help the rubber polymer chains get back to their original form after stretching. (d) Used: to manufacture tyres, rubber hose and gloves Vulcanised rubber Unvulcanised rubber Can withstand high temperatures Cannot withstand high temperatures More elastic Less elastic Harder and stronger Less hard, softer Can resist oxidation Easily oxidised 4.3 Synthetic Rubber 1. Synthetic rubber is man-made polymer produced from polymerisation process of hydrocarbon obtained from fractions of petroleum.
80SOS TIP 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 Styrene-Butadiene Rubber (SBR) Styrene and butadiene, C4 H6 • Resistant to high heat • Easily vulcanised • Corrosion resistance Tyres Shoe soles Neoprene Chloroprene, C4 H5 Cl • Elastic • Resistant to high heat • Strong tensile strength • Resistant to oxidation • Not inflammable Water pipe, gloves, electrical wire insulators, petrol rubber hose Butyl rubber Isobutylene, C4 H8 • Less permeable than natural rubber • Resistant to heat • Can be vulcanised with sulphur under normal condition Inner tube of tyres, roofing materials, nonpermeable layer to prevent diffusion of water in reservoir Thiokol rubber 1,2-dichloroethane, C2 H4 Cl2 Very resistant against chemical solvents and oils Lining for oils and solvents storage tanks 4. Advantages of synthetic rubber: (a) Withstand inclement weather and high temperatures. (b) Non-inflammable and can withstand heat. (c) Good heat and electrical insulators. (d) Not easily permeable by gas and water. (e) Very high resistance towards chemicals. 5. Disadvantages of synthetic rubber: (a) Ability to absorb vibrations, sound, shock is less compared to natural rubber. (b) Elasticity and tensile strength are less compared to natural rubber. Each question is followed by four options A, B, C or D . Choose the best option for each question. PAPER 1 4.1 Polymers 1. Which of the following is the correct pair of natural polymer and its monomer? Natural polymer Monomer A Polyisoprene Propene B Starch Amino acid C Cellulose Glucose D Protein Isoprene 2. Which substance is a natural polymer? A Polythene C Polyisoprene B Polypropene D Polyvinyl chloride 3. Which polymer has the empirical formula of CH2 ? A H H C H C C n 6 H5 C H H C H C C n 2 H5 B H H C H C H n D H H C H C CH n 3 Question 2: Polyisoprene is the name of natural rubber. Question 3: Empirical formula for polymers is the same as its monomer.
87SOS TIP Section A Answer all questions. 1. The following flow chart shows the change of latex to vulcanised rubber through Steps I and II. S2 Cl2 Acid Step I Step II Solid natural rubber Latex Vulcanised rubber 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] PAPER 2 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.
93SOS 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 Initial length (cm) 8.0 cm 8.0 cm Length with 200 g weight (cm) 0 2 3 4 5 6 7 8 9 10 11 12 7 8 9 10 11 12 Clip Rubber P Weight 200 g Clip Rubber Q Weight 200 g 0 2 3 4 5 6 7 8 9 10 11 12 Length after weight is withdrawn (cm) 10.0 cm 8.0 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: • Vulcanisations produce sulphur cross linkages between rubber polymer chains. • Vulcanisations produce rubber that is more elastic and of better quality.
114 CHAPTER 1 Paper 1 1. D Acidified potassium dichromate(VI) is an oxidising agent. It oxidises alcohols to carboxylic acids. 2. B Potassium bromide is a reducing agent. Bromide ion is oxidised to bromine. 3. A (+3) + 3(x) = 0 3 + 3x = 0 x = −1 4. B Carbon can reduce zinc oxide but cannot reduce aluminium oxide. 5. C KMnO4 solution acts as an oxidising agent under acidic condition. Sulphuric acid supplies the H+ ions to KMnO4 . 6. A Combustion of metal in oxygen produces metal oxides are redox reactions. 7. C A displacement reaction occurred; shows that copper is more reactive than metal M. It is more likely to release electrons and be oxidised. 8. B Iodide ion is a reducing agent; easily oxidised to iodine. 9. B Combustion and rusting are redox reactions 10. B SiO2 + C ⎯→ Si + CO2 +4 0 11. D Reduction involves loss of oxygen, gain of electrons or gain of hydrogen. 12. B Chlorine water is an oxidising agent that oxidises Fe2+ to Fe3+. 13. C Magnesium displaces Cu2+ ions from copper(II) nitrate. Excess magnesium powder also observed in the beaker. 14. D 1(+1) + (Cl) + 4(−2) = 0 Cl = +7 15. B A reducing agent releases electrons. 16. C Zinc displaces Cu2+ from the solution. When all Cu2+ ions are displaced, blue colour disappears and Cu2+ ions change to Cu atoms, producing a brown solid. 17. A Al and Ca are above Ag in the reactivity series of metals. Sulphur is a non-metal. 18. B Chlorine, Cl2 is an oxidising agent; oxidises I− to I2 . Iodine solution is brown. 19. D Acidified KMnO4 is an oxidising agent. SO2 and H2 S are reducing agents that will react with KMnO4 , CH4 is hydrocarbon. 20. B The chemical formula of chromium(VI) oxide is CrO3 . 21. D A more reactive metal reduces the oxide of a less reactive metal. Q > R > S > P 22. B Graphite and platinum are inert electrodes. 23. D The more positive the E° value, the easier reduction occurs at the cathode (positive terminal). E0 cell = (+0.34) – (−0.44) = + 0.78 V 24. B The more negative the E° value, the more readily for oxidation to occur; stronger reduction agent. 25. D The more positive the E0 value, the more readily for reduction to occur; stronger oxidation agent. 26. A E0 value of Zn is more negative than the E0 value of Fe. Fe is the cathode and Zn is the anode. E0 cell = –0.44 – (–0.76) = +0.32 V 27. A The right-hand side is a chemical cell, S (−) is anode and R (+) is cathode. The left-hand side is an electrolytic cell: P is cathode (−) and Q is anode (+). Reduction occurs at cathodes, R and P. 28. D The right-hand side is a chemical cell: Mg is anode (−); Mg is oxidised to Mg2+. The left-hand side is an electrolytic cell: Cu is anode (+); Cu is oxidised to Cu2+. 29. B Cell I: Oxidation of Cu to Cu2+. Cell II: Oxidation of Cl− to Cl2 . 30. C Anode (P) is plating metal, cathode (Q) is object to be plated and electrolyte (R) is silver nitrate. 31. D E0 of OH– is more negative. OH– is oxidised to O2 gas. E0 of OH+ is more positive, H+ is reduced to H2 gas. 32. B Higher concentration of Cl– , thus Cl– ions easier to be oxidised at the anode. 33. B Cells I & III: Hydrogen gas is produced at the cathode and chlorine gas is produced at the anode. Cell II: Zinc formed at the anode, clorine gas formed at the cathode. 34. C Anode: O2 gas; 4OH− → O2 + 4e− + 2H2 O Cathode: H2 gas; 2H+ + 2e− → H2 35. C Hydrogen gas is produced at the cathode and bromine gas is produced at the anode. Na+ ions combine with OH− ions to produce sodium hydroxide, NaOH. 36. B Cell notation of a voltaic cell is written as: Electrode(s) ǀ Electrolyte(aq) ǁ Electrolyte(aq) ǀ Electrode(s) Electrode Ionic solution Ionic solution Electrode Anode (negative terminal) Cathode (positive terminal) 37. B Cu2+ ions are reduced at the cathode to deposit copper metal onto the electrode. Cu2+(aq) + 2e– ⎯→ Cu(s) 38. A Bauxite is an aluminium ore. Aluminium is extracted by electrolysis of molten bauxite. 39. A Metals above carbon in the reactivity series of metals are extracted by electrolysis of their molten ores. Carbon is positioned between aluminium and zinc. 40. C Carbon reduces oxides of zinc and copper. Zinc reduces oxides of iron and copper. Electrolysis is used in purification of metal. 41. C Ca is above Mg in the reactivity series of metals. 42. D Iron(III) oxide is reduced by carbon and carbon monoxide to iron and carbon dioxide. 43. D Al is formed at the cathode by reduction. Al3+ ions gain electrons to form Al atoms. 44. C Carbon (coke) is more reactive than iron but less reactive than aluminium. Carbon (coke) reduces iron oxide but not aluminium oxide. 45. A Cryolite has a low melting point than bauxite. It reduces the high melting point of bauxite thus save energy. 46. D A sacrificial metal must be more electropositive than iron. A sacrificial metal lose its electrons easier. Sacrificial metal ionises before iron, corrosion of iron is prevented. Chapter 1 Answers
115 47. B Copper is the least electropositive metal that has the slowest rate of corrosion. 48. B Magnesium is more electropositive than iron. Manesium release electrons more easily than iron. Magnesium undergoes oxidation. Rusting of iron is prevented. 49. A Aluminium has a protective oxide layer against corrosion. Aluminium and aluminium oxide are amphoteric. 50. D Sodium is a reactive metal and reacts vigorously with air and water. Paper 2 Section A 1. (a) Electrolysis is a process to decompose a compound into its elements by passing an electric current through the compound in its molten or aqueous state. (b) Acid X: Hydrochloric acid, HCl(aq) Acid Y: Nitric acid, HNO3 (aq) (c) Chorine gas (d) (i) 2H+(aq) + 2e− → H2 (g) (ii) 2H2 O(l) → O2 (g) + 4H+(aq) + 4e− (e) (i) Use a lower concentration of hydrochloric acid. The higher the concentration of Cl ions, the easier for Cl− ion to be oxidised to chlorine gas. (ii) 2Cl−(aq) → Cl2 (g) + 2e− 2. (a) Electrolyte is a compound that conducts an electric current in the molten or aqueous state and undergo chemical changes. (b) (i) There is a flow of current/electrons in the circuit. (ii) Copper(II) ion, Cu2+ and chloride ion, Cl− . (iii) Yellow-green gas is liberated. (iv) 2Cl−(l) → Cl2 (g) + 2e− (c) (i) The electrode connected to the negative terminal of the battery. (ii) Brown solid deposited on the cathode. (iii) Cu2+(aq) + 2e− → Cu(s) 3. (a) (i) A redox reaction involves oxidation and reduction taking place simultaneously. (ii) C(s) + 2PbO(s) → 2Pb(s) + CO2 (g) (iii) +2 to 0 (iv) Lead(II) oxide (b) Carbon is more reactive than metal X. Carbon is less reactive than zinc. (c) (i) Zinc, carbon, lead, X (ii) X is copper (d) Crucible Metal oxide + carbon 4. (a) Zn(s) + Cu2+ (aq) → Zn2+(aq) + Cu(s) (b) (i) and (ii) Voltmeter Cathode Copper electrode Electrode 2 Anode Copper(II) sulphate solution (c) Voltmeter reading (V) Metal 1.10 Zinc 0.78 Iron 0.38 Tin 0.00 Copper (d) Magnesium 5. (a) Electrolysis of water (b) Sodium hydroxide (c) (i) and (ii) Anode: Oxidation (loss of electrons) Cathode: Reduction (gain of electrons) (d) The product is water, a non-toxic substance 6. (a) Haematite (b) (i) Limestone, CaCO3 (ii) To remove calcium silicate (slag) through reaction between silicon dioxide and calcium oxide. CaO + SiO2 → CaSiO3 (slag) (c) (i) Carbon monoxide (ii) C + CO2 → 2CO 7. (a) (i) Fe2+ (ii) Fe3+ (b) (i) P: Magnesium/Aluminium/Zinc [any one] (ii) Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) (c) (i) Q: Carbon (ii) 3C(s) + 2Fe2 O3 (s) → 3CO2 (g) + 4Fe(s) Section B 8. (a) Oxidising agent: Hydrogen peroxide Reducing agent: Iodide ion (b) Oxidation: 2I−(aq) → I2 (aq) + 2e− Reduction: H2 O2 (aq) + 2H+(aq) + 2e− → 2H2 O(l) (c) Iodine: −1 to 0 Oxygen: −1 to -2 Hydrogen: no change, +1 (d) Colourless solution turns brown (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 Chapter 1
116 (ii) Set I: Iron(II) ion, Fe2+ Set II: Iodide ion, I− (iii) Set I: Fe2+(aq) → Fe3+(aq) + e− Set II: 2I−(aq) → I2 (aq) + 2e− Section C 10. (a) Iron undergoes oxidation at the anode: Fe(s)→ Fe2+(aq) + 2e− Electrons flow through iron to the cathode (edge of water droplet). Oxygen is reduced to hydroxide ion, OH− . O2 + 2H2 O + 4e− → 4OH Fe2+ ions combined with OH− ions to form a green precipitate, iron(II) hydroxide. Fe2+(aq) + 2OH−(aq) → Fe(OH)2 (s) Further oxidation by oxygen, produced brown precipitate, hydrated iron(III) oxide or rust. Fe(OH2 )(s) ⎯⎯⎯→ Fe2 O3 .xH2 O(s) oxidation (b) (i) Metal X: Magnesium Metal Y: Lead Magnesium is more electropositive than iron. Magnesium is more likely to release electrons compared to iron and oxidised. Mg(s) → Mg2+(aq) + 2e− Magnesium preventing iron from rusting. While lead is less electropositive than iron. Iron releases electrons and oxidised. Iron will rusts. Fe(s) → Fe2+(aq) + 2e− (ii) Procedure: Test tube I Iron nail coiled with zinc strip Iron nail coiled with tin strip Iron nail Iron nail coiled with aluminium strip Iron nail coiled with copper strip Agar solution + potassium hexacyanoferrate(III) + phenolphthalein indicator Test tube II Test tube IV Test tube V Test tube III 1. Iron nails together with metal strips are cleaned using sandpaper. 2. Four iron nails are each coiled with different metals. 3. The five iron nails are placed into separate test tubes containing agar solution mixed with potassium hexacyanoferrate(III) and phenolphtalein. 4. The test tubes are left in a test tube rack for two days. 5. All changes are recorded. Observation: Test tube Dark blue spot Intensity of pink colour Inference I None High Iron nail did not rust II Most Low Iron nail rusted the most III None High Iron nail did not rust IV Moderate Low Iron nail rusted a bit V Few - Iron nail rusted a bit Discussion: Potassium hexacyanoferrate(III) solution confirms the presence of Fe2+ ions which are produced when iron rusts. The more the dark blue spots formed, the higher the corrosion of iron. Phenolphthalein detects the presence of OH− ions which are produced at the cathode during the corrosion of metals. The higher the intensity of the pink colour, the greater the extent of rusting of iron. A metal more electropositive than iron can protect the iron from rusting. A less electropositive metal than iron speeds up the corrosion of iron. 11. (a) (i) Solution X in Set I: Dilute hydrochloric acid At the anode, the E0 value of OH– is less positive that the E0 value of Cl– . Oxidation of water to O2 gas occurs more readily. Oxygen gas is colourless. Solution X in Set II: Concentrated hydrochloric acid At the anode, the concentration of Cl− ion is higher than OH– ion. Oxidation of Cl− to Cl2 gas occurs more readily. Chlorine gas is yellow-green. (ii) H+ ions and Cl− ions (iii) Set I: Anode: 4OH−(aq) → O2 (g) + 2H2O(l) + 4e− Cathode: 2H+(aq) + 2e− → H2 (g) Set II: Anode: 2Cl−(aq) → Cl2 (g) + 2e− Cathode: 2H+(aq) + 2e− → H2 (g) (b) Metal P: Iron (Fe); Metal Q: Copper (Cu); Metal R: Zinc (Zn) P: Iron(II) nitrate; Q: Copper(II) nitrate; R: Zinc nitrate (i) Materials: Zinc foil, iron foil, iron(II) nitrate solution, copper(II) nitrate solution Apparatus: Beaker (ii) Procedure: Zinc foil Beaker I II III Iron foil Iron(II) nitrate solution Copper(II) nitrate solution 1. Set up the apparatus as shown above with a metal dipped into a nitrate salt solution. 2. Observe the changes that occur in each beaker. Results: Beaker I: Zinc dissolves; grey solid deposited; green solution turns colourless. Beaker II: Zinc dissolves; brown solid deposited; blue solution turns colourless. Beaker III: Iron dissolves; brown solid deposited; blue solution turns colourless. Conclusion: Zinc is the most electropositive metal. It displaces iron and copper from their nitrate salt solutions. Iron is more electropositive than copper. It displaces copper from its salt solution. Cu, Fe, Zn more reactive Chapter 1