Suggested Teaching Scheme

Transcription

1 Suggested Teaching Scheme Suggested Teaching Scheme The following suggested teaching schemes are for teachers reference only. Teachers may revise them based on the time-tabling arrangement of their own schools. Scheme 1: Chemistry to be studied in Secondary 3, 4, 5 and 6 In many schools, the Chemistry curriculum is studied in Secondary 3, 4, 5 and 6. Although the distribution of periods varies from school to school, the total number of periods for the curriculum is generally around 416. A possible distribution of periods is as follows: A possible distribution of periods for S3, S4, S5 and S6 S3 S4 S5 S6 Number of teaching weeks per year Number of periods per week Total number of periods per year Total number of periods for the curriculum 416 Suggested teaching scheme for the curriculum Level S3 (56 periods) S4 (140 periods) S5 (140 periods) S6 (80 periods) Only 2 out of 3 Content Suggested number of period(s) Topic 1 Planet Earth 12 Topic 2 Microscopic World I 44 Revision on laboratory safety 1 Topic 3 Metals 39 Topic 4 Acids and Bases 45 Topic 5 Redox Reactions, Chemical Cells and Electrolysis 41 Topic 6 Microscopic World II 14 Revision on laboratory safety 1 Topic 7 Fossil Fuels and Carbon Compounds 32 Topic 8 Chemistry of Carbon Compounds 45 Topic 9 Chemical Reactions and Energy 13 Topic 10 Rate of Reaction 16 Topic 11 Chemical Equilibrium 18 Topic 12 Patterns in the Chemical World 15 Revision on laboratory safety 1 Topic 13 Industrial Chemistry 39 Topic 14 Materials Chemistry 39 Topic 15 Analytical Chemistry 40 Schools taking investigative study need to allocate an extra of 30 periods for the curriculum.

2 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Scheme 2: Chemistry to be studied in Secondary 4, 5 and 6 In some schools, the Chemistry curriculum is studied in Secondary 4, 5 and 6. The total number of periods for the curriculum is generally around 360. A possible distribution of periods is as follows: A possible distribution of periods for S4, S5 and S6 S4 S5 S6 Number of teaching weeks per year Number of periods per week Total number of periods per year Total number of periods for the curriculum 360 Suggested teaching scheme for the curriculum Level S4 (140 periods) S5 (140 periods) S6 (80 periods) Only 2 out of 3 Content Suggested number of period(s) Topic 1 Planet Earth 8 Topic 2 Microscopic World I 31 Topic 3 Metals 32 Topic 4 Acids and Bases 36 Topic 5 Redox Reactions, Chemical Cells and Electrolysis 33 Revision on laboratory safety 1 Topic 6 Microscopic World II 13 Topic 7 Fossil Fuels and Carbon Compounds 29 Topic 8 Chemistry of Carbon Compounds 41 Topic 9 Chemical Reactions and Energy 12 Topic 10 Rate of Reaction 15 Topic 11 Chemical Equilibrium 16 Topic 12 Patterns in the Chemical World 13 Revision on laboratory safety 1 Topic 13 Industrial Chemistry 39 Topic 14 Materials Chemistry 39 Topic 15 Analytical Chemistry 40 Schools taking investigative study need to allocate an extra of 30 periods for the curriculum. 2

3 Suggested Teaching Scheme Suggested number of periods for Topic 5 Chemistry for Total number of periods Suggested number of periods for each unit S3 S6 (Scheme 1) 41 Unit 18 Chemical cells in daily life Unit 19 Simple chemical cells Unit 20 Oxidation and reduction Unit 21 Oxidation and reduction in chemical cells Unit 22 Electrolysis S4 S6 (Scheme 2) 33 Unit 18 Chemical cells in daily life Unit 19 Simple chemical cells Unit 20 Oxidation and reduction Unit 21 Oxidation and reduction in chemical cells Unit 22 Electrolysis

4 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Teaching Plan An extension of the study of the reactivity of metals in Topic 3 leads to the study of electron flow in the external circuit when two different metals are immersed in an electrolyte. This is a chemical cell, a device in which chemical energy is converted into electrical energy. Students will study the characteristics of common primary cells and secondary cells in Unit 18. Reactions in simple chemical cells and the electrochemical series of metals are discussed in Unit 19. To enhance students' understanding of the chemistry involved in a chemical cell, the concept of redox reactions is introduced in Unit 20. Students will carry out investigations involving common oxidizing and reducing agents. They will also learn how to write chemical equations for redox reactions. With the concepts related to redox reactions, students will study reactions occurring in more complicated chemical cells (such as lead-acid accumulator and fuel cell) in Unit 21. Unit 22 introduces electrolysis, the use of electricity to bring out chemical reactions. Students will study reactions that occur during electrolysis. In addition, students should be able to predict products in electrolysis according to the different factors affecting the preferential discharge of ions. The concepts of redox reactions have a number of applications in industrial chemistry and daily life. Through searching for information and critically reading articles about electrochemical technology, students should appreciate the contribution of chemical knowledge to technological innovations, which in turn improve our quality of life. Students should also be able to assess the environmental impact and safety issues associated with these technologies. Organization of the topic Redox Reactions, Chemical Cells and Electrolysis Unit 18 Chemical cells in daily life Unit 19 Simple chemical cells Unit 20 Oxidation and reduction Unit 21 Oxidation and reduction in chemical cells Unit 22 Electrolysis 4

5 Teaching Plan Unit 18 Chemical cells in daily life Section Key point(s) Suggested task(s) for students Remark Total number of period = Electricity from chemical reactions 18.2 Different types of chemical cells 18.3 Terms related to chemical cells Total number of period = 1 What a chemical cell is Magnesium-copper chemical cell Primary cells and secondary cells Battery Negative and positive electrodes Electrolyte Cell capacity Discharge Service life, cycle life and shelf life 18.4 Zinc-carbon cell Zinc casing as the negative electrode Carbon rod as the positive electrode Ammonium chloride as the electrolyte Refer to the following website for an animation illustrating how chemical cells power a flashlight: iastate.edu/greenbowe/ sections/projectfolder/ animations/flashlightv8. html (accessed July 2014) 18.5 Alkaline manganese cell Zinc as the negative electrode Manganese(IV) oxide as the positive electrode Potassium hydroxide as the electrolyte Continued on next page 5

6 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Section Key point(s) 18.6 Silver oxide cell Zinc as the negative electrode Silver oxide as the positive electrode Potassium hydroxide as the electrolyte Total number of period = Lithium ion cell Lithium atoms lying between graphite sheets as the negative electrode Lithium metal oxide as the positive electrode Lithium salt as the electrolyte 18.8 Nickel metal hydride (NiMH) cell Hydrogen absorbing alloys as the negative electrode Nickel(II) hydroxide as the positive electrode Potassium hydroxide as the electrolyte Suggested task(s) for students Remark Application manuals of different types of cells from a manufacturer com/static. aspx?name= AppManuals An online cell guidebook radioshack.com/ support_tutorials/ batteries/batgd-b.htm (accessed July 2014) Continued on next page 6

7 Teaching Plan Section 18.9 Lead-acid accumulator Total number of period = Choosing a chemical cell for a particular use Environmental impact of using chemical cells Key point(s) Lead plates as the negative electrode Lead plates coated with lead(iv) oxide as the positive electrode Sulphuric acid as the electrolyte Aspects to consider when choosing a chemical cell for a particular use Heavy metal components of chemical cells are toxic Advantages of using secondary cells Suggested task(s) for students Decision Making Remark 7

8 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Unit 19 Simple chemical cells Section Key point(s) Suggested task(s) for students Remark Total number of period = Reactions in simple chemical cells How electrons flow in the external circuit of a simple chemical cell Ionic half-equation Activity 19.1 Building simple chemical cells Refer to the following websites for animations illustrating the chemical reactions occurring in a zinc-copper chemical cell: iastate.edu/ Greenbowe/sections/ projectfolder/ animations/cuzncell. html com/physsci/ chemistry/ essentialchemistry/ flash/galvan5.swf The following website allows students to build computer simulated chemical cells using given materials: iastate.edu/greenbowe/ sections/projectfolder/ flashfiles/electrochem/ volticcell.html (accessed July 2014) Continued on next page 8

9 Teaching Plan Section 19.2 Lemon cells made from different metal couples Key point(s) Measuring the voltages of lemon cells formed when coupling other metals with copper Comparing the tendency of metals to form ions based on voltages of the lemon cells Suggested task(s) for students Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) 19.3 The electrochemical series of metals 19.4 Improving simple chemical cells 19.5 The role of a salt bridge Introducing the electrochemical series of metals Comparing the electrochemical series with the reactivity series Separating a simple chemical cell into two half-cells Functions of a salt bridge Practice 19.1 Activity 19.2 Determining the order of three metals in the electrochemical series Practice 19.2 Remark Chemical cell with a salt bridge Total number of periods = The Daniell cell Structure of a Daniell cell Reactions occurring in a Daniell cell Refer to the following website for an animation illustrating the chemical reactions occurring in a Daniell cell: chemistry_interactive/ daniell_cell.htm (accessed July 2014) 9

10 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Unit 20 Oxidation and reduction Section Key point(s) Suggested task(s) for students Remark Total number of period = Defining oxidation and reduction in terms of gain and loss of oxygen Describing oxidation and reduction in terms of gain and loss of oxygen Redox reaction 20.2 Defining oxidation and reduction in terms of loss and gain of hydrogen Total number of period = Defining oxidation and reduction in terms of loss and gain of electrons 20.4 Oxidizing agent and reducing agent Total number of period = Relative strength of reducing and oxidizing agents Total number of period = 1 Describing oxidation and reduction in terms of loss and gain of hydrogen Describing oxidation and reduction in terms of electron transfer Introducing oxidizing agent and reducing agent Defining oxidizing and reducing agents in terms of electron transfer The trend of reducing power of metals in the electrochemical series The trend of oxidizing power of metal ions in the electrochemical series 20.6 Oxidation numbers The concept of oxidation number Rules for assigning oxidation numbers to elements in different species Practice 20.1 Practice 20.2 Continued on next page 10

11 Teaching Plan Section Total number of periods = Defining oxidation and reduction in terms of changes in oxidation numbers 20.8 Using oxidation numbers to identify the oxidizing agent and reducing agent in a redox reaction 20.9 Advantages and disadvantages of using the concept of oxidation number The Stock system of naming compounds Total number of periods = Common oxidizing and reducing agents Key point(s) Using changes in oxidation number to identify redox reactions Using changes in oxidation number to identify the oxidizing agent and reducing agent in a redox reaction Advantages and disadvantages of using the oxidation number concept Naming cations Naming polyatomic anions Chemical changes of common oxidizing and reducing agents Ionic half-equations representing the chemical changes Suggested task(s) for students Practice 20.3 Activity 20.1 Investigating redox reactions Practice 20.4 Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) Balancing redox equations using ionic half-equations Balancing redox equations using oxidation number method Total number of period = The electrochemical series and the relative oxidizing / reducing power of common oxidizing / reducing agents How to balance redox equations using ionic half-equations How to balance redox equations using oxidation number method Introducing a detailed version of the electrochemical series Predicting the feasibility of a redox reaction using the electrochemical series Practice 20.5 Remark Continued on next page 11

12 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Section Key point(s) Suggested task(s) for students Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) Chlorine as an oxidizing agent Action of aqueous chlorine on potassium bromide solution Action of aqueous chlorine on potassium iodide solution Reaction with sodium hydroxide solution Activity 20.2 Ranking halogens according to their oxidizing power Practice 20.6 Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) Nitric acid of different concentrations as oxidizing agents Oxidizing property of concentrated and dilute nitric acids Activity 20.3 Investigating the action of nitric acid of different concentrations on metals Practice 20.7 Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) Concentrated sulphuric acid as an oxidizing agent Action of concentrated sulphuric acid on metals Action of concentrated sulphuric acid on nonmetals Action of concentrated sulphuric acid on halides Activity 20.4 Investigating the properties of concentrated sulphuric acid Practice 20.8 Remark Reaction of aqueous chlorine with halide solutions Comparing the action of nitric acid of different concentrations on magnesium and copper Action of concentrated sulphuric acid on zinc and copper Action of concentrated sulphuric acid on halides Total number of period = Aqueous sulphur dioxide as a reducing agent Action of aqueous sulphur dioxide on some common oxidizing agents Test for sulphur dioxide gas Practice 20.9 Investigating the chemical properties of sulphur dioxide (I) (microscale experiment) Investigating the chemical properties of sulphur dioxide (II) (microscale experiment) 12

13 Teaching Plan Unit 21 Oxidation and reduction in chemical cells Section Key point(s) Suggested task(s) for students Remark Total number of periods = Oxidation and reduction in a simple chemical cell Reactions occurring at the electrodes of a chemical cell Practice Redox reactions in a zinc-carbon cell Redox reactions in a zinc-carbon cell Two main disadvantages of the cell Refer to the following website for an animation illustrating the chemical reactions occurring in a zinccarbon cell: chem.iastate.edu/ Greenbowe/sections/ projectfolder/animations/ ZnCbatteryV8web.html (accessed July 2014) Total number of period = Redox reactions in simple chemical cells with inert electrodes Redox reactions in simple chemical cells set up using acidified K 2 Cr 2 O 7 (aq) and FeSO 4 (aq) KI(aq) and Br 2 (aq) Practice 21.2 Redox reactions in simple chemical cells (using carbon electrodes) Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) 21.4 Fuel cells How a hydrogen-oxygen fuel cell works Advantages and limitations of hydrogenoxygen fuel cells Decision making Lithium or hydrogen powered vehicles Chemistry magazine Oxygen absorbers for packaged foods Fuel cell Refer to the following websites for animations illustrating how a fuel cell works: orgs/mpa/mpa11/ animation.htm org/energy/ commercial/print/cell_ is_print_e.html (accessed July 2014) 13

14 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Unit 22 Electrolysis Section Key point(s) Suggested task(s) for students Remark Total number of period = Electrolysis: chemical reactions from electricity 22.2 Comparing a chemical cell and an electrolytic cell 22.3 Electrolysis of molten sodium chloride using carbon electrodes 22.4 Some knowledge related to aqueous electrolytes Total number of periods = Electrolysis of aqueous solutions of ionic compounds Terms commonly used in electrolysis Function Direction of electron flow Reactions at electrodes Illustrating the chemical changes brought out by electricity using electrolysis of molten sodium chloride as an example Dissociation of water Dissociation of acids in water Dissolving electrolytes in water Electrolysis of acidified water using platinum electrodes Electrolysis of very dilute sodium chloride solution using carbon electrodes Activity 22.1 Investigating the electrolysis of acidified water Practice 22.1 Electrolysis of acidified water Refer to the following website for an animation illustrating the chemical reactions occurring during the electrolysis of acidified water: (accessed July 2014) Electrolysis of very dilute sodium chloride solution using carbon electrodes Continued on next page 14

15 Teaching Plan Section Key point(s) Suggested task(s) for students Remark Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) 22.6 Factors affecting the order of discharge of ions during the electrolysis of aqueous solutions 22.7 The position of ions in the electrochemical series and the order of discharge of ions The position of ions in the electrochemical series The effect of concentration of ions in the solution The nature of the electrodes Order of discharge of cations Order of discharge of anions Activity 22.2 Investigating factors affecting the order of discharge of ions during electrolysis position of ions in the electrochemical series Total number of periods = 2 (Scheme 1), total number of period = 1 (Scheme 2) 22.8 The effect of concentration of ions in the solution and the order of discharge of ions Electrolysis of dilute or concentrated sodium chloride solution using carbon electrodes Activity 22.3 Investigating factors affecting the order of discharge of ions during electrolysis effect of concentration of ions in the solution Electrolysis of concentrated sodium chloride solution using carbon electrodes Practive 22.2 Total number of periods = The nature of electrodes and the order of discharge of ions Electrolysis of dilute copper(ii) sulphate solution using carbon electrodes Electrolysis of dilute copper(ii) sulphate solution using copper electrodes Electrolysis of concentrated sodium chloride solution using a mercury cathode Activity 22.4 Investigating factors affecting the order of discharge of ions during electrolysis effect of the nature of electrodes Practice 22.3 Practice 22.4 Manufacture of chlorine by electrolysis of brine Mercury electrolytic cell Continued on next page 15

16 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Section Key point(s) Suggested task(s) for students Remark Total number of period = Industrial uses of electrolysis Refining of copper Electroplating Activity 22.5 Electroplating with nickel Practice 22.5 Electroplating with nickel Electroplating Environmental impact of the electroplating industry Pollutions due to acids, alkalis, compounds of heavy metals and cyanides Methods to control pollution from the electroplating industry 16

17 Teaching Notes Teaching Notes Unit 20 Oxidation and reduction N2 page 42 Oxidation number of sulphur in different substances Examination questions often ask about the oxidation number of sulphur in different substances. Substance Oxidation number of S in the substance H 2 S 2 Na 2 S 2 O 3 +2 SO 2 +4 NaHSO 3 +4 Na 2 SO 4 +6 H 2 S 2 O 7 +6 N3 page 42 Oxidation number of nitrogen in different substances Examination questions often ask about the oxidation number of nitrogen in different substances. Substance Oxidation number of N in the substance NH 3 3 N 2 O +1 NO +2 HNO 2 +3 NO 2 +4 HNO

18 Topic 5 Redox Reactions, Chemical Cells and Electrolysis N4 page 44 Using oxidation numbers to identify unfamiliar redox reactions Examination questions often list equations of unfamiliar reactions and ask students to determine whether the reactions involve oxidation and reduction. Students may use changes in oxidation numbers to identify the redox reactions. Examples : Pb(s) + PbO 2 (s) + 2H 2 SO 4 (aq) 2PbSO 4 (s) + 2H 2 O(l) H 2 O 2 (aq) + H 2 SO 4 (aq) + 2KI(aq) K 2 SO 4 (aq) + I 2 (aq) + 2H 2 O(l) NH 4 NO 3 (s) N 2 O(g) + 2H 2 O(l) Fe 2 (SO 4 ) 3 (aq) + H 2 S(g) 2FeSO 4 (aq) + S(s) + H 2 SO 4 (aq) KClO 3 (s) 2KCl(s) + 3O 2 (g) FeSO 4 (s) Fe 2 O 3 (s) + SO 3 (g) + SO 2 (g) MnO 2 (s) + 4HCl(aq) MnCl 2 (aq) + Cl 2 (g) + 2H 2 O(l) N5 page 46 Using oxidation numbers to identify the oxidizing agent and reducing agent in an unfamiliar redox reaction Examination questions often list equations of unfamiliar reactions and ask students to identify the oxidizing agents (i.e. species being reduced) or the reducing agents (i.e. species being oxidized). Examples: +4 0 SO 2 + 2Mg 2MgO + S SO 2 is the oxidizing agent, it is being reduced +4 0 SO 2 + 2H 2 S 3S + 2H 2 O SO 2 is the oxidizing agent, it is being reduced 3 0 2NH 3 + 3CuO 3Cu + N 2 + 3H 2 O NH 3 is the reducing agent, it is being oxidized 0 +2 Zn + 2AgNO 3 Zn(NO 3 ) 2 + 2Ag Zn is the reducing agent, it is being oxidized +2 0 CuSO 4 + Zn ZnSO4 + Cu CuSO 4 is the oxidizing agent, it is being reduced Fe 2 O 3 + 3CO 2Fe + 3CO2 CO is the reducing agent, it is being oxidized 1 0 Fe 2 (SO 4 ) 3 + 2KI 2FeSO 4 + K 2 SO 4 + I 2 KI is the reducing agent, it is being oxidized 18

19 Teaching Notes N7 page 51 Distinguishing between Fe 2+ (aq) ions and Fe 3+ (aq) ions Examination questions often ask about methods to distinguish between Fe 2+ (aq) ions and Fe 3+ (aq) ions: observing their colours; treating with acidified potassium permanganate solution; treating with dilute sodium hydroxide solution / dilute aqueous ammonia; treating with concentrated nitric acid. Colour Fe 2+ (aq) ions are pale green in colour while Fe 3+ (aq) ions are yellow-brown in colour. Acidified potassium permanganate solution Fe 2+ (aq) ions can decolorize acidified potassium permanganate solution but Fe 3+ (aq) ions cannot. MnO 4 (aq) + 5Fe 2+ (aq) + 8H + (aq) Mn 2+ (aq) + 5Fe 3+ (aq) + 4H 2 O(l) Dilute sodium hydroxide solution / dilute aqueous ammonia Fe 2+ (aq) ions give a green precipitate with dilute sodium hydroxide solution / dilute aqueous ammonia. Fe 3+ (aq) ions give a reddish brown precipitate with dilute sodium hydroxide solution / dilute aqueous ammonia. Fe 2+ (aq) + 2OH (aq) Fe 3+ (aq) + 3OH (aq) Fe(OH) 2 (s) Fe(OH) 3 (s) Concentrated nitric acid Fe 2+ (aq) ions give a brown gas with concentrated nitric acid but Fe 3+ (aq) ions do not. Refer to N16 for further details. 19

20 Topic 5 Redox Reactions, Chemical Cells and Electrolysis N8 page 59 Oxidizing / reducing action of hydrogen peroxide Examination questions may ask about the oxidizing / reducing action of hydrogen peroxide, an unfamiliar reagent. Students need to write the ionic half-equations for the chemical changes based on the given reactants and products. Hydrogen peroxide can act as an oxidizing agent: H 2 O 2 (aq) + 2H + (aq) + 2e 2H 2 O(l) and as a reducing agent: H 2 O 2 (aq) O 2 (g) + 2H + (aq) + 2e Acidic solutions favour the oxidizing action. e.g. oxidizing Fe 2+ (aq) ions to Fe 3+ (aq) ions H 2 O 2 (aq) + 2Fe 2+ (aq) + 2H + (aq) 2Fe 3+ (aq) + 2H 2 O(l) oxidizing I (aq) ions to I 2 (aq) H 2 O 2 (aq) + 2I (aq) + 2H + (aq) I 2 (aq) + 2H 2 O(l) Alkaline solutions favour the reducing action. e.g. reducing Cl 2 (aq) to Cl (aq) ions H 2 O 2 (aq) + Cl 2 (aq) O 2 (g) + 2Cl (aq) + 2H + (aq) N15 page 66 Similarity in chemical properties of halogens Examination questions aften ask about the similarity in chemical properties of halogens. Examples: Both Cl 2 (aq) and Br 2 (aq) can oxidize SO 2 3 (aq) ions to SO 2 4 (aq) ions. X 2 (aq) + SO 3 2 (aq) + H 2 O(l) 2X (aq) + SO 4 2 (aq) + 2H + (aq) Both Cl 2 (aq) and Br 2 (aq) can react with Fe 2+ (aq) ions to give Fe 3+ (aq) ions. X 2 (aq) + 2Fe 2+ (aq) 2X (aq) + 2Fe 3+ (aq) Both Cl 2 (aq) and Br 2 (aq) can react with I (aq) ions to give I 2 (aq). X 2 (aq) + 2I (aq) 2X (aq) + I 2 (aq) Both Cl 2 (aq) and Br 2 (aq) can undergo disproportionation in alkalis. X 2 (g) + 2OH (aq) X (aq) + OX (aq) + H 2 O(l) Refer to N22 for the gradual change in chemical properties of halogens / halide ions. 20

21 Teaching Notes N16 page 67 Action of concentrated nitric acid on metals, iron(ll) salts and sulphites 1 Metals Concentrated nitric acid oxidizes most metals, for example, magnesium and copper. However, it has no reaction with iron and aluminium. Concentrated nitric acid renders them completely passive. This is due to the formation of a layer of oxide on the metal surface. 2 Iron(II) salts Concentrated nitric acid oxidizes iron(ll) salts to iron(lll) salts. Nitrogen monoxide is produced and this reacts with oxygen in air to form brown nitrogen dioxide gas. 3Fe 2+ (aq) + NO 3 (aq) + 4H + (aq) 3Fe 3+ (aq) + NO(g) + 2H 2 O(l) 2NO(g) + O 2 (g) 2NO 2 (g) 3 Sulphites Concentrated nitric acid oxidizes sulphites to sulphates. SO 3 2 (aq) + 2H + (aq) + 2NO 3 (aq) SO 4 2 (aq) + 2NO 2 (g) + H 2 O(l) 21

22 Topic 5 Redox Reactions, Chemical Cells and Electrolysis N17 page 68 Comparing the properties of dilute nitric acid and dilute sulphuric acid / dilute hydrochloric acid Examination questions often ask students to compare the properties of different acids. Comparing the properties of dilute HNO 3 (aq) and dilute H 2 SO 4 (aq) Acid Property Observation Dilute HNO 3 (aq) Dilute H 2 SO 4 (aq) Reaction with alkali Reaction with carbonate (or hydrogencarbonate) Action on litmus solution salt and water are produced carbon dioxide gas is given off give a red colour Reaction with zinc a brown gas is given off a colourless gas is given off Reaction with barium chloride solution no white precipitate forms a white precipitate forms Titration with NaOH(aq) more NaOH(aq) is needed to reach the end point for H 2 SO 4 (aq) than HNO 3 (aq) Comparing the properties of dilute HNO 3 (aq) and dilute HCl(aq) Property Acid Observation Dilute HNO 3 (aq) Dilute HCl(aq) Reaction with alkali Reaction with carbonate (or hydrogencarbonate) Action on litmus solution salt and water are produced carbon dioxide gas is given off give a red colour Reaction with copper a brown gas is given off no gas is given off Reaction with silver nitrate solution no white precipitate forms a white precipitate forms 22

23 Teaching Notes N18 page 69 Comparing the properties of concentrated sulphuric acid and dilute sulphuric acid Examination questions often ask students to compare the properties of different acids. Comparing the properties of concentrated H 2 SO 4 (l) and dilute H 2 SO 4 (aq) Property Acid Observation Concentrated H 2 SO 4 (l) Dilute H 2 SO 4 (aq) Reaction with alkali Reaction with carbonate (or hydrogencarbonate) Action on litmus solution salt and water are produced carbon dioxide gas is given off give a red colour Oxidizing property can oxidize copper no reaction with copper Non-volatility non-volatile, can displace other volatile acids (e.g. HCl(g) and HNO 3 (g)) from their salts not non-volatile N20 page 69 Preparation of sulphur dioxide in the laboratory In the laboratory, we can prepare sulphur dioxide by heating copper turnings with concentrated sulphuric acid. The gas is dried by passing through concentrated sulphuric acid. The gas is collected by downward delivery because it is denser than air. Cu(s) + 2H 2 SO 4 (l) CuSO 4 (aq) + SO 2 (g) + 2H 2 O(l) 23

24 Topic 5 Redox Reactions, Chemical Cells and Electrolysis N21 page 69 Distinguishing between concentrated nitric acid and concentrated sulphuric acid Copper can be used to distinguish between concentrated nitric acid and concentrated sulphuric acid. Copper gives a brown gas (NO 2 ) with concentrated nitric acid. Copper gives a colourless gas (SO 2 ) with concentrated sulphuric acid. N22 page 69 Gradual change in chemical properties of halogens / halide ions Examination questions often ask about the gradual change in chemical properties of halogens / halide ions. Examples: order of increasing oxidizing power: I 2 < Br 2 < Cl 2 Aqueous chlorine can displace bromine from potassium bromide solution and displace iodine from potassium iodide solution. Aqueous bromine can displace iodine from potassium iodide solution but cannot displace chlorine from potassium chloride solution. order of increasing reducing power: I > Br > Cl NaCl(s) is not oxidized by concentrated H 2 SO 4 (l); NaBr(s) reacts with concentrated H 2 SO 4 (l) to give Br 2 (l) and SO 2 (g); NaI(s) reacts with concentrated H 2 SO 4 (l) to give I 2 (s) and H 2 S(g). 24

25 Teaching Notes Unit 21 Oxidation and reduction in chemical cells N1 page 85 Redox reactions occurring in unfamiliar chemical cells Examination questions often show unfamiliar chemical cells and ask students to write ionic half-equations for redox reactions occurring in the cells based on their knowledge on oxidation and reduction. Examples: A chemical cell made up of an aluminium can, a carbon rod and household bleach At the aluminium can The aluminium can undergoes oxidation to give tetrahydroxylaluminate ions, [Al(OH) 4 ] (aq). Al(s) + 4OH (aq) [Al(OH) 4 ] (aq) + 3e At the carbon rod The hypochlorite ions, OCl (aq) (in the household bleach) undergo reduction to give chloride ions and hydroxide ions. OCl (aq) + H 2 O(l) + 2e Cl (aq) + 2OH (aq) A chemical cell made up of copper electrodes and copper(ll) sulphate solution of different concentrations (concentration cell) 25

26 Topic 5 Redox Reactions, Chemical Cells and Electrolysis In the set-up, electrons flow in such a direction that the concentration of Cu 2+ (aq) ions in each half-cell becomes the same eventually, i.e. electrons flow from X to Y in the external circuit. At electrode X Cu(s) Cu 2+ (aq) + 2e At electrode Y Cu 2+ (aq) + 2e Cu(s) A sodium-sulphur cell operating at about 370 ºC (above the melting points of sodium and sulphur) At electrode A Na(l) Na + (l) + e At electrode B S(l) + 2e S 2 (l) N4 page 89 Predicting chemical changes occurring in a chemical cell based on the oxidizing power of the reagents involved Examination questions may ask students to predict the chemical changes that would occur in a chemical cell based on the oxidizing power of the reagents involved. Example: consider the following chemical cell: 26

27 Teaching Notes 1 To tackle the question, first identify the oxidizing agent and the reducing agent. The question gives the information that Br 2 (aq) is a stronger oxidizing agent than Fe 3+ (aq) ion. Hence it can be deduced that Br 2 (aq) acts as the oxidizing agent and Fe 2+ (aq) ion acts as the reducing agent. The following chemical changes would occur: Fe 2+ (aq) Fe 3+ (aq) + e Br 2 (aq) + 2e 2Br (aq) 2 Based on the chemical changes above, it is possible to identify the negative electrode (anode) and the positive electrode (cathode) of the chemical cell. As oxidation occurs at electrode X, thus it is the negative electrode, i.e. the anode. As reduction occurs at electrode Y, thus it is the positive electrode, i.e. the cathode. 3 Based on the above information, it is possible to decide that the direction of electron flow in the external circuit is from electrode X to electrode Y. N5 page 89 A simple chemical cell involving KI(aq) and Fe 2 (SO 4 ) 3 (aq) Fe 2 (SO 4 ) 3 (aq) is a strong oxidizing agent than I 2 (aq). 27

28 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Hence the following changes occur in the chemical cell: 2I (aq) I 2 (aq) + 2e Fe 3+ (aq) + e Fe 2+ (aq) If some starch solution is added to the concentrated potassium iodide solution, a blue colour would appear after some time. This is because iodine forms a complex with starch. Unit 22 Electrolysis N8 page 116 Electrolysis of dilute copper(ii) sulphate solution using carbon anode and copper cathode Besides using carbon electrodes / copper electrodes, examination questions may ask about the electrolysis of dilute copper(ll) sulphate solution using carbon anode and copper cathode. At the anode 4OH (aq) O 2 (g) + 2H 2 O(l) + 4e At the cathode Cu 2+ (aq) + 2e Cu(s) The blue colour of the solution fades gradually because the concentration of copper(ii) ions in the electrolyte decreases. Copper(II) ions and hydroxide ions are consumed in the electrolysis. Hydrogen ions and sulphate ions remain in the solution. Thus the solution eventually becomes sulphuric acid. 28

29 Suggested Answers Suggested Answers page 1 1 Rusting refers to the corrosion of iron ONLY. 2 Fe 2 O 3 xh 2 O(s) 3 Reactivity series Unit 18 Chemical cells in daily life Decision Making page 12 There is no right or wrong answer to this question. Students may choose according to one of the following criteria: Choose alkaline manganese cell in every case for the sake of convenience. This saves the trouble of replacing dry cells frequently and recharging of cells. Choose based on the cost-effectiveness of the different types of cell. Choose rechargeable nickel metal hydride cell if dry cells are used for a long time every day. Operating cost Type of cell Torch (HKD per hour) Camera flash unit (HKD per time) Portable CD player (HKD per hour) Radio (HKD per hour) Zinc-carbon Alkaline manganese It is more cost effective to use alkaline manganese cell for torch, camera flash unit and portable CD player, while it is more cost effective to use zinc-carbon cell for radio. Unit Exercise pages a) electrical b) electrolyte c) primary cell d) zinc-carbon cell e) alkaline manganese cell f) silver oxide cell 29

30 Topic 5 Redox Reactions, Chemical Cells and Electrolysis g) lithium ion cell h) nickel metal hydride cell i) lead-acid accumulator 2 a) voltmeter; digital multimeter b) battery c) negative; positive d) electrolyte e) service f) shelf 3 Chemical cell Rechargeable or not? Material of negative electrode Material of positive electrode Material of electrolyte Maximum voltage (V) Zinc-carbon cell No zinc carbon ammonium chloride 1.5 Alkaline manganese cell No zinc manganese(iv) oxide potassium hydroxide 1.5 Silver oxide cell No zinc silver oxide potassium hydroxide 1.5 Lithium ion cell Yes lithium atoms lying between graphite sheets lithium metal oxide lithium salt dissolved in an organic solvent 3.7 Nickel metal hydride cell Yes hydrogenabsorbing alloy nickel(ii) hydroxide potassium hydroxide 1.2 Lead-acid accumulator Yes lead titanium plates coated with lead(iv) oxide sulphuric acid 2 4 A Option B No electrons flows in the external circuit as both electrodes are made of copper. Options C and D No electrons flows in the external circuit as distilled water and ethanol do NOT conduct electricity at all. 5 B 6 C 30

31 Suggested Answers 7 A 8 C Option C Nickel metal hydride cells have a relatively high rate of self-discharge. 9 D 10 D 11 B (1) The electrolyte of a zinc-carbon cell is ammonium chloride. (3) A zinc-carbon cell shows poor performance in high-drained devices. 12 B (2) The maximum voltage of both alkaline manganese cell and zinc-carbon cell is 1.5 V. 13 B 14 D 15 A (2) The electrolyte of a nickel metal hydride cell is potassium hydroxide. (3) A nickel metal hydride cell shows good performance in high-drained devices. 16 a) In any aircraft or spacecraft. b) Any one of the following: Life-support systems required during a power failure, in a remote area or in an ambulance. Medical aids that cannot run from a normal power supply, such as heart pacemakers and bionic ears. Other sensitive and medical instrumentation used in a remote area. c) Any one of the following: Inside any computer or other electronic system. Inside the human body. 17 a) Anode lithium atoms lying between graphite sheets Cathode lithium metal oxide b) A lithium salt dissolved in an organic solvent c) light weight high energy density 31

32 Topic 5 Redox Reactions, Chemical Cells and Electrolysis (d) Any one of the following: Electric razors Electric toothbrushes Medical equipment Portable DVD players PDAs Laptop computers 18 Answers for the HKALE question are not provided. 19 a) 1.5 V b) Any one of the following: Use alkaline manganese cells reasons these cells have a steady voltage during discharge; there is no need to recharge the used cells. Use nickel metal hydride cells reasons more cost effective as these cells can be recharged over 500 times; the physical volume of discarded cells in landfills can be reduced. Unit 19 Simple chemical cells Practice P19.1 page 24 a) Zinc b) Form zinc to copper c) i) Cu 2+ (aq) + 2e Cu(s) ii) Zn(s) Zn 2+ (aq) + 2e d) The voltage of the cell becomes zero. This is because sugar solution contains no mobile ion and thus does not allow ionic conduction between the two electrodes. 32

33 Suggested Answers P19.2 page 26 a) Metal X b) i) P is salt bridge. A salt bridge serves two important functions: It completes the circuit by allowing ions to move from one half-cell to the other. It provides ions that can move into the half-cells to prevent the build-up of charge in the solutions which would cause the reaction to stop. ii) Metal X iii) From metal X to metal Y Unit Exercise pages a) electrolyte b) positive c) electrons d) negative e) positive 2 a) electrons; zinc ions; Zn(s) Zn 2+ (aq) + 2e b) electrons; silver; Ag + (s) + e Ag(s) c) zinc; silver d) negative; positive 3 C 4 D Zinc forms ions more readily than silver does. Thus, electrons flow from zinc to silver in the external circuit. 5 D Option A The silver strip acts as the positive electrode, i.e. the cathode. Option B The magnesium strip acts as the negative electrode, i.e. the anode. Option C Silver ions in beaker X gain electrons and form silver atoms. Ag + (aq) + e Ag(s) Thus, the concentration of silver ions in beaker X decreases. 33

34 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Option D There is a build up of negative charge in beaker X due to the NO 3 (aq) ions that remained. K + (aq) ions migrate from the salt bridge into beaker X to offset any buildup of negative charge. K + ions from the salt bridge electron flow NO 3 ions from the salt bridge e e Ag+ Mg 2+ 6 B 7 B Option A The zinc strip acts as the negative electrode, i.e. the anode. Option B Electrons flow from the zinc strip to the copper container in the external circuit. Option C The porous pot does NOT provide ions for charge balance. Option D Copper(II) ions gain electrons and form copper atoms. The concentration of copper(ii) ions in the copper(ii) sulphate solution decreases. Thus, the blue colour of the solution fades. 8 A (3) Zinc atoms lose electrons and form zinc ions. Zn(s) Zn 2+ (aq) + 2e This does not affect the colour of the dilute sulphuric acid. 9 C (2) Electrons move from the zinc plate to the copper plate in the external circuit. (3) The difference between the tendencies for magnesium and copper to form ions is greater than between zinc and copper. Thus, the voltage of the cell would increase if the zinc plate is replaced by a magnesium plate. 34

35 Suggested Answers 10 A (1) Magnesium is more reactive than nickel. Thus, magnesium atoms will lose electrons and form magnesium ions. Mg(s) Mg 2+ (aq) + 2e Thus, concentration of magnesium ions in beaker X increases. (2) Nickel(II) ions in the nickel(ii) sulphate solution gain electrons and form nickel atoms. Ni 2+ (aq) + 2e Ni(s) Thus, the mass of the nickel electrode increases gradually. (3) The salt bridge provides ions that can move into the half-cells to prevent the build-up of charge in the solutions. 11 a) Fe b) Ag c) Iron and silver. They are furthest apart in the electrochemical series. 12 a) Cadmium b) Cd(s) + Ni 2+ (aq) Cd 2+ (aq) + Ni(s) c) From the cadmium strip to the nickel strip d) When the cell operates, Cd 2+ (aq) ions will form at the cadmium strip. NO 3 (aq) ions in the salt bridge will migrate into cadmium half-cell to offset any buildup of Cd 2+ (aq) ions. 13 a) It completes the circuit by allowing ions to move from one half-cell to the other. It provides ions that can move into the half-cells to prevent the build-up of charge in the solutions which would cause the reaction to stop. b) iron < X < Y Metal X is more reactive than iron. Thus, X forms ions more readily than iron does. Metal Y is more reactive than metal X. Thus, Y forms ions more readily than X does. c) i) Fe 2+ (aq) + 2e Fe(s) X(s) X 2+ (aq) + 2e ii) The green colour of solution fades out, because concentration of Fe 2+ decreases. iii) Voltage increased, because copper forms ions less readily than iron does. 35

36 Topic 5 Redox Reactions, Chemical Cells and Electrolysis 14 a) b) Electrons flow from the zinc rod to the copper can in the external circuit. c) No current would flow through the conducting wires. This was because the glass beaker would not allow ions to move between the two solutions. d) The reading on the ammeter would decrease. 15 a) Electrons flow from sodium and electrode B via external circuit towards electrode A and nickel(ii), because sodium is more reactive than nickel. b) Na(l) Na + (l) + e c) To melt the electrolyte and mobilize the ions. d) This cell provides a high voltage. 16 Answers for the HKCEE question are not provided. Unit 20 Oxidation and reduction Practice P20.1 page 41 1 a) An oxidation is involved as Zn(s) loses electrons. b) A reduction is involved as Fe 3+ (aq) gains electrons. c) An oxidation is involved as I (aq) loses electrons. 36

37 Suggested Answers 2 Cu(s) Ag + (aq) Ionic half-equation for chemical change that occurs for the species Cu(s) Cu 2+ (aq) + 2e Ag + (aq) + e Ag(s) Whether the species undergoes oxidation or reduction? oxidation reduction P20.2 page 43 a) Oxidation number of O = 2 Suppose the oxidation number of N in NO is x. x + ( 2) = 0 x = +2 \ the oxidation number of N in NO is +2. b) Oxidation number of H = +1 Suppose the oxidation number of N in NH 3 is x. x + (+1) x3 = 0 x = 3 \ the oxidation number of N in NH 3 is 3. c) Oxidation number of Na = +1 Oxidation number of O = 2 Suppose the oxidation number of C in Na 2 CO 3 is x. [(+1) x 2 + x + ( 2) x 3] = 0 x = +4 \ the oxidation number of C in Na 2 CO 3 is +4. d) Cu(OH) 2 consists of Cu 2+ ion and OH ions. Oxidation number of Cu = charge on the ion = +2 e) Oxidation number of Fe = charge on the ion = +3 f) Oxidation number of O = 2 Suppose the oxidation number of Cr in Cr 2 O 7 2 is x. [(x) x 2 + ( 2) x 7] = 2 x = +6 \ the oxidation number of Cr in Cr 2 O 7 2 is

38 Topic 5 Redox Reactions, Chemical Cells and Electrolysis P20.3 page 47 1 a) from +5 decrease to +4 b) from 0 increase to +2 c) from +4 decrease to +2 2 a) Yes. The oxidation number of Al increases from 0 to +3, so Al(s) is reducing agent. The oxidation number of H decreases from +1 to 0, so H + (aq) is oxidizing agent. b) No. The oxidation number of Cr, O, H remain +6, 2 and +1 respectively. c) Yes. The oxidation number of Mg increases from 0 to +2, so Mg(s) is reducing agent. The oxidation number of Cu decreases from +2 to 0, so Cu 2+ (aq) is oxidizing agent. P20.4 page 52 Acidified potassium dichromate solution Iron(II) sulphate solution Product formed when it reacts Chromium(III) Iron(III) Ionic half-equation As an oxidizing agent or a reducing agent? Cr 2 O 7 2 (aq) + 14H + (aq) + 6e oxidizing agent 2Cr 3+ (aq) + 7H 2 O(l) Fe 2+ (aq) Fe 3+ (aq) + e reducing agent P20.5 page Zn(s) + 14H + (aq) + Cr 2 O 2 7 (aq) 3Zn 2+ (aq) + 7H 2 O(l) + 2Cr 3+ (aq) 2 a) Reduction: MnO 4 (aq) + 8H + (aq) + 5e Mn 2+ (aq) + 4H 2 O(l) Oxidation: SO 2 3 (aq) + H 2 O(l) SO 2 4 (aq) + 2H + (aq) + 2e b) 5SO 2 3 (aq) + 6H + (aq) + 2MnO 4 (aq) 5SO 2 4 (aq) + 3H 2 O(l) + 2Mn 2+ (aq) c) Reducing agent: SO 2 3 (aq) Oxidizing agent: MnO 4 (aq) d) Purple permanganate solution is becomes colourless. 38

39 Suggested Answers P20.6 page 66 Add the sample solution to bromine solution. Potassium iodide solution changes the solution brown, but potassium chloride solution does not change. P20.7 page 68 Acid Action of acid on magnesium Name of gas given off, if any Action of acid on copper Name of gas given off, if any Property shown by the acid (acidic / oxidizing property) Very dilute nitric acid yes, hydrogen no, acidic property Dilute nitric acid yes, nitrogen monoxide yes, nitrogen monoxide oxidizing property Concentrated nitric acid yes, nitrogen dioxide yes, nitrogen dioxide oxidizing property P20.8 page 71 Dilute sulphuric acid Concentrated sulphuric acid Action of acid on Name of gas given off, if any Property shown by the acid (acidic / oxidizing property) Name of gas given off, if any Property shown by the acid (acidic / oxidizing property) Sodium carbonate carbon dioxide acidic property carbon dioxide acidic property Zinc hydrogen acidic property sulphur dioxide oxidizing property Copper no reaction acidic property sulphur dioxide oxidizing property P20.9 page 73 Add bromine solution / iodine solution / acidified potassium permanganate / acidified potassium dichromate Sodium sulphite solution changes the bromine solution / iodine solution / acidified potassium permanganate colourless, but sodium sulphate solution does not change. Sodium sulphite solution changes acidified potassium dichromate green, but sodium sulphate solution does not change. 39

40 Topic 5 Redox Reactions, Chemical Cells and Electrolysis Unit Exercise pages a) reduction b) oxidation c) oxidation d) oxidation e) reducing 2 Species that causes Species that is Reducing agent oxidation / reduction oxidized / reduced Oxidizing agent oxidation / reduction oxidized / reduced 3 oxidizing power of metal ion increasing weak oxidizing agents strong oxidizing agents Metal ion K + (aq) + e Ca 2+ (aq) + 2e Na + (aq) + e Mg 2+ (aq) + 2e Al 3+ (aq) + 3e Zn 2+ (aq) + 2e Fe 2+ (aq) + 2e Pb 2+ (aq) + 2e 2H + (aq) + 2e Cu 2+ (aq) + 2e Ag + (aq) + e Au + (aq) + e Metal K(s) Ca(s) Na(s) Mg(s) Al(s) Zn(s) Fe(s) Pb(s) H 2 (g) Cu(s) Ag(s) Au(s) strong reducing agents weak reducing agents reducing power of metal decreasing 4 a) Oxidizing agent Product when reduced Ionic half-equation Aqueous chlorine Cl (aq) Cl 2 (g) + 2e 2Cl (aq) Aqueous bromine Br (aq) Br 2 (aq) + 2e 2Br (aq) Potassium permanganate in acidic solution Mn 2+ (aq) MnO 4 (aq) + 8H + (aq) + 5e Mn 2+ (aq) + 4H 2 O(l) Potassium dichromate in acidic solution Cr 2 O 7 2 (aq) 2Cr 3+ (aq) Cr 2 O 2 7 (aq) + 14H + (aq) + 6e 2Cr 3+ (aq) + 7H 2 O(l) Iron(III) sulphate solution Fe 3+ (aq) Fe 2+ (aq) Fe 3+ (aq) + e Fe 2+ (aq) 40

41 Suggested Answers b) Oxidizing agent Product when oxidized Ionic half-equation Sodium sulphite solution SO 4 2 (aq) SO 3 2 (aq) + H 2 O(l) SO 4 2 (aq) + 2H + (aq) + 2e Iron(II) sulphate solution Fe 2+ (aq) Fe 3+ (aq) Fe 2+ (aq) Fe 3+ (aq) + e Potassium bromide solution 2Br (aq) Br 2 (aq) 2Br (aq) Br 2 (aq) + 2e Potassium iodide solution 2l (aq) l 2 (aq) 2l (aq) l 2 (aq) + 2e 5 B 6 A 7 C 8 C 9 C 10 A 11 B 12 A 13 A 14 B 15 D 16 A 17 D 18 a) Oxidation number of O = 2 Suppose the oxidation number of S in SO 3 is x. x + ( 2) x 3 = 0 x = +6 the oxidation number of S in SO 3 is

42 Topic 5 Redox Reactions, Chemical Cells and Electrolysis b) Oxidation number of O = 2 Suppose the oxidation number of Mn in MnO 4 is x. x + ( 2) x 4 = 1 x = +7 the oxidation number of Mn in MnO 4 is +7. c) Oxidation number of H = +1 Oxidation number of O = 2 Suppose the oxidation number of C in HCO 3 is x. (+1) + x + ( 2) x 3 = 1 x = +4 the oxidation number of C in HCO 3 is +4. d) Oxidation number of H = +1 Oxidation number of O = 2 Suppose the oxidation number of S in H 2 S 2 O 7 is x. (+1) x 2 + 2x + ( 2) x 7 = 0 2x = +12 x = +6 the oxidation number of S in H 2 S 2 O 7 is +6. e) PbSO 4 consists of Pb 2+ ion and SO 4 2 ion. Oxidation number of Pb = charge on the ion = +2 f) Oxidation number of Ca = +2 Suppose the oxidation number of H in CaH 2 is x. (+2) + (x) x 2 = 0 x = 1 the oxidation number of H in CaH 2 is 1. g) Oxidation number of H = +1 Oxidation number of O = 2 Suppose the oxidation number of Al in [Al(OH) 4 ] is x. x + [( 2) + (+1) x 4] = 1 x = +3 the oxidation number of Al in [Al(OH) 4 ] is +3. h) Oxidation number of H = +1 Suppose the oxidation number of O in H 2 O 2 is x. (+1) x 2 + (x) x 2 = 0 x = 1 the oxidation number of O in H 2 O 2 is -1. i) Oxidation number of Cl = 1 Suppose the oxidation number of Co in [CoCl 4 ] 2 is x. x + [( 1) x 4] = 2 x = +2 the oxidation number of Co in [CoCl 4 ] 2 is

43 Suggested Answers j) Oxidation number of O = 2 Suppose the oxidation number of V in VO 2 + is x. x + 2x( 2) = +1 x = +5 the oxidation number of V in VO 2 + is a) This is not a redox reaction because the oxidation numbers of all elements remain unchanged in the reaction NaI(s) + H 2 SO 4 (l) NaHSO 4 (s) + HI(g) b) The oxidation number of K increases from 0 to +1 while that of H decreases from +1 to 0. Therefore it is a redox reaction K(s) + 2H 2 O(l) 2KOH(aq) + H 2 (g) The oxidation number of H decreases, i.e. H 2 O is being reduced. c) The oxidation number of N increases from 3 to 0 while that of Cu decreases from +2 to 0. Therefore it is a redox reaction NH 3 (g) + 3CuO(s) 3Cu(s) + N 2 (g) + 3H 2 O(g) The oxidation number of Cu decreases, i.e. CuO is being reduced. d) The oxidation number of Zn increases from 0 to +2 while that of Ag decreases from +1 to 0. Therefore it is a redox reaction Zn(s) + 2AgNO 3 (aq) Zn(NO 3 ) 2 (s) + 2Ag(s) The oxidation number of Ag decreases, i.e. AgNO 3 is being reduced. e) The oxidation number of Br increases from 1 to 0 while that of Cl decreases from 0 to 1. Therefore it is a redox reaction KBr(aq) + Cl 2 (aq) 2KCl(aq) + Br 2 (aq) The oxidation number of Cl decreases, i.e. Cl 2 is being reduced. 20 a) 2Cr 2 O 7 2 (aq) + 6I (aq) + 14H + (aq) 2Cr 3+ (aq) + 3I 2 (s) + 7H 2 O(l) b) 2MnO 4 (aq) + 10Cl (aq) + 16H + (aq) 2Mn 2+ (aq) + 5Cl 2 (g) + 8H 2 O(l) c) 2Fe 2+ (aq) + H 2 O 2 (aq) + 2H + (aq) 2Fe 3+ (aq) + 2H 2 O(l) 43