Oxidation and Reduction

1. Oxidation and Reduction Mr Field

2. Using this slide show • The slide show is here to provide structure to the lessons, but not to limit them….go off-piste when you need to! • Slide shows should be shared with students (preferable electronic to save paper) and they should add their own notes as they go along. • A good tip for students to improve understanding of the calculations is to get them to highlight numbers in the question and through the maths in different colours so they can see where numbers are coming from and going to. • The slide show is designed for my teaching style, and contains only the bare minimum of explanation, which I will elaborate on as I present it. Please adapt it to your teaching style, and add any notes that you feel necessary.

3. Menu: • Lesson 1 – Oxidation and Reduction • Lesson 2 – Redox Equations • Lesson 3 – Reactivity Series • Lesson 4 – Voltaic Cells • Lesson 5-6 – Electrolytic Cells • Lesson 7 – HL – Standard Electrode Potentials • Lesson 8 – HL – Ecell and Non-Standard Conditions • Lesson 9 – HL – Advanced Electrolysis • Lesson 10 – HL – Quantitative Electrolysis • Lesson 11–12 – Internal Assessment

4. Lesson 1 Oxidation and Reduction

5. Overview • Copy this onto an A4 page. You should add to it as a regular review throughout the unit.

6. Assessment • This unit will be assessed by: • An internal assessment at the end of the topic • A joint test at the end of the Acids and Bases topic

7. We Are Here

8. Lesson 1: Oxidation and Reduction • Objectives: • Reflect on prior knowledge of oxidation and reduction • Understand oxidation and reduction in terms of electron transfer • Calculate oxidation numbers

9. Reflecting on Redox • Write down everything you know on oxidation and reduction • You have 60 seconds

10. Defining Oxidation and Reduction • O - oxidation • I - is • L - loss of electrons • R - reduction • I - is • G - gain of electrons • Often (but far from always) in practice: • Oxidation is gain of oxygen or loss of hydrogen • This results in the loss of electrons • Reduction is loss of oxygen or gain of hydrogen • This results in the gain of electrons

11. Oxidation States / Oxidation Numbers • Oxidation state is the charge an atom would have if all it’s bonds were ionic • It is the number of electrons an atom has gained or lost by forming bonds • You even talk about oxidation state of covalent compounds! • It is important as the oxidation state of an atom has a significant impact on its chemistry • Fe(II) Fe(III) • Cr(III) Cr(VI) • Mn(II) Mn(VII)

12. Calculating Oxidation States • The oxidation state of an element is zero • The oxidation states of a neutral compound sum to zero, and of an ion sum to the charge on the ion • The more electronegative atom in an ion assumes a negative oxidation state, the less electronegative one a positive oxidation state • Some rules of thumb: • Start with these, work the others out.

13. For example • Determine the oxidation states of each atom in the following: • CO2 • O,-2 -2 except with F or a peroxide • C, +4 to balance out the 2 lots of ‘-2’ • H2SO4 • O, -2 -2 except with F or peroxide • H, +1 +1 except in metal hydrides • S, +6 +6 since four lots of ‘-2’ and two of ‘-1’ sum to -6 • BaO2 (barium peroxide) • Ba, +2 +2 since Gp II metal • O, -1 -1 since peroxide • CO32- • O, -2 -2 except with F or peroxide • C, +4 +4 since 3 x -2 and +4 sums to the charge, -2 Note: oxidation states must be written with the sign in front: +2 NOT 2+

14. Determine oxidation states for each atom in: • H2O • S8 • CH4 • H3PO4 • CCl4 • HClO • KMnO4 • IO3- • Cr2O72- • Cr(H2O)63+

15. Oxidation States and Names • Oxidation states are used in the names of compounds • ‘ate’ means an element is in a positive oxidation state • Usually because it is bonded with oxygen • ‘ide’ means an element is in a negative oxidation state • Oxidation state of transition metals is given in Roman numerals • FeCl2– iron (II) chloride • Iron (II) means Fe in the +2 ox. State • Chloride means the chlorine is in a negative oxidation state • FeCl3– iron (III) chloride • Iron (III) mean Fe in the +3 ox. State • KClO3– potassium chlorate • Chlorate tells you the chlorine is in a positive oxidation state

16. Name the following: • MnO2 • CuO • Cu2O • KMnO4 • K2Cr2O7

17. Redox Reactions • Whenever an oxidation occurs, a reduction also occurs, hence REDOX • For example: Zn(s) + Cu2+(aq)  Cu(s) + Zn2+(aq) 0 +2 0 +2 • Zinc is oxidised – it loses two electrons • Zinc is the reducing agent because it reduces copper • Copper is reduced – it gains two electrons • Cu2+ is the oxidising agent because it oxidises the zinc • The number of electrons gained by species is always equal to the number of electrons lost by species.

18. Disproportionation* • When some atoms of an element are oxidised and others are reduced • For example: 2 H2O2 2H2O + O2 +1 -1 +1 -2 0 • The oxygen ending in the H2O loses an electron and is oxidised • The oxygen ending in the O2 gains an electron and is reduced *This is not on the syllabus but is useful and interesting

19. Identify the species that are oxidised and reduced in each reaction, stating the number of electrons gained or lost and stating whether disproportionation takes place • Fe2O3 + 2 Al  Al2O3 + 2 Fe • AgNO3 + NaCl  AgCl + NaNO3 • 3 Cl2 + 6 OH− → 5 Cl− + ClO3− + 3 H2O • H2SO4 + 2HBr  Br2 + SO2 + 2H2O • Cu + 4 HNO3  Cu(NO3)2 + 2 NO2 + 2 H2O

20. Key Points • Oxidation state tells us the number of electorns an atom has gained or lost • The most electronegative atom in a bond gains electrons, to form a negative oxidation state and vice versa • Oxidation reactions are always accompanied by reductions

21. Lesson 2 Redox Equations

22. Refresh • Fertilizers may cause health problems for babies because nitrates can change into nitrites in water used for drinking. • Define oxidation in terms of oxidation numbers. • Deduce the oxidation states of nitrogen in the nitrate, NO3–, and nitrite, NO2–, ions.

23. We Are Here

24. Lesson 2: Redox Equations • Objectives: • Deduce simple half-equations • Combine half-equations to form full equations • Conduct a series of redox reactions • Use H+ and H2O to balance redox equations

25. Half-equations • Half equations show the changes to individual species in a redox reaction. • Fe2O3 + 2 Al  2 Fe + Al2O3 • Fe3+ + 3 e-  Fe ….this is the reduction • Al  Al3+ + 3 e- ….this is the oxidation • A wide variety of half equations can be found in the data booklet

26. Combining Half Equations and Balancing Redox Reactions – Example 1

27. Combining Half Equations and Balancing Redox Reactions – Example 2

28. Microscale Redox Reactions • Complete the experiment here, on microscale redox reactions • Rather than doing it on a plastic sheet, use a dropping tile • For each change you observe, you should write a balanced redox equation to describe it • Once you finish, you should practice balancing the redox equations on the following slide

29. Produce balanced redox equations for the reactions of: • Bromine with sodium • ½ Br2 + e- Br- • Na Na+ + e- • Copper (II) oxide with hydrogen • Cu2+ + 2e-  Cu • ½ H2  H+ + e- • Aluminium reacting with chromate ions • Al  Al3+ + 3e- • Cr2O72- + 6e- 2Cr3+ • Iron (II) chloride reacting with manganate ions • Fe2+  Fe3+ + e- • MnO4- + 5 e- Mn2+

30. Key Points • Half-equations show the changes to each species in a redox reaction • To combine half-equations into a full equation • Multiply each such that the electron-transfers balance • Add H2O to balance O • Add H+ to balance H • Add e- to balance charge

31. Lesson 3 Reactivity Series

32. Refresh • Nitric acid reacts with silver in a redox reaction: __ Ag(s) + __ NO3–(aq) + ____ → __ Ag+(aq) + __ NO(g) + ____ • Using oxidation numbers, deduce the complete balanced equation for the reaction showing all the reactants and products.

33. We Are Here

34. Lesson 3: Reactivity Series • Objectives: • Deduce reactivity series from chemical observations • Use reactivity series to predict the feasibility of reactions

35. Redox and Reactivity • Redox behaviour is closely linked to reactivity • The most reactive metals are the best reducing agents • The most reactive non-metals are the best oxidising agents • The least reactive elements are neither good oxidising or reducing agents

36. Reactivity and Displacement Reactions • Reactive metals are better reducing agents than un-reactive metals. • As such, a reactive metal can displace less reactive metals from their compounds. Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • As a good reducing agent, the zinc reduces the Cu2+, causing it to gain two electrons.

37. Constructing a Reactivity Series • Complete the experiment here in which you have to construct a reactivity series. • Analysis • Use your reactivity series to predict the feasibility of the reactions of: • I- with Fe3+ • Zn with Sn2+ • Fe2+with Cl • Compare your series with the order found on Table 14 of the data booklet.

38. Key Points • More reactive metals are better reducing agents • More reactive non-metals with oxidising agents • A more reactive metal will reduce (displace) ions of a less reactive metal

39. Lesson 4 Voltaic Cells

40. Refresh • Consider the following three redox reactions. Cd(s) + Ni2+(aq) → Cd2+(aq) + Ni(s) Ni(s) + 2Ag+(aq) → Ni2+(aq) + 2Ag(s) Zn(s) + Cd2+(aq) → Zn2+(aq) + Cd(s) • Deduce the order of reactivity of the four metals, cadmium, nickel, silver and zinc and list in order of decreasing reactivity. • Identify the best oxidizing agent and the best reducing agent.

41. We Are Here

42. Lesson 4: Voltaic Cells • Objectives: • Explain in simple terms how voltaic cells use redox reactions to produce electricity • Understand that oxidation occurs at the anode and reduction at the cathode • Make a series of voltaic cells in order to better understand the how they work

43. Voltaic Cells • The reaction of Mg with Cu2+ ions: • Mg(s) + Cu2+(aq) Mg2+(aq) + Cu(s) • This reaction involves two electrons being transferred from the Mg to the Cu: • Mg  Mg2+ + 2e- • Cu2+ + 2e-  Cu • The Mg reduces the copper ions as it is more reactive • This is an exothermic reaction, and the energy is normally released as heat • A voltaic cell forces each half of the reaction to take place in a separate container, with the electrons moving through a circuit to get from one side to the next • This is an exothermic reaction, where the energy is released as electrical rather than thermal energy • The reactions in Voltaic cells usually involve only metals but do not have to.

44. Voltaic Cells Continued Anode: Where oxidation happens Electron Flow - + Electron Flow Cathode: Where reduction happens

45. Key Parts of a Voltaic Cell • Anode • Electrode or ‘half-cell’ where oxidation happens • Contains the more reactive metal • The negative electrode: produces electrons • Cathode • Electrode or ‘half-cell’ where reduction happens • Contains the less reactive metal • The positive electrode: accepts electrons • Salt Bridge • Contains a neutral salt such as potassium nitrate • Made of a tube of jelly or a filter paper soaked in salt solution • Ions diffuse in and out to balance charge and complete circuit • Voltmeter • Measures the difference in potential between half-cells • Could be replaced with other circuitry to do useful work • REMEMBER • AnOx • Anode-Oxidation • CaRe • Cathode-Reduction

46. Drawing a cell • Draw and fully label a zinc/iron cell. Include: • Labels for cathode and anode • Labels for positive and negative • Each half-equation • Arrow showing direction of electron flow

47. Constructing Voltaic Cells • You will need to build and measure the potential of voltaic cells comprising various combinations of the following: • Cu/Cu2+ • Fe/Fe2+ • Mg/Mg2+ • Sn/Sn2+ • Zn/Zn2+ • Follow the instructions here

48. Key Points • Voltaic cells extract electrical energy from redox reactions by separating each half • At the anode, the more reactive of two metals is oxidised • At the cathode, the less reactive of two metals is reduced

49. Lesson 5-6 Electrolytic Cells

50. Refresh A particular voltaic cell is made from magnesium and iron half-cells. The overall equation for the reaction occurring in the cell is Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s) Which statement is correct when the cell produces electricity? • Magnesium atoms lose electrons. • The mass of the iron electrode decreases. • Electrons flow from the iron half-cell to the magnesium half-cell. • Negative ions flow through the salt bridge from the magnesium half-cell to the iron half-cell. • For each incorrect statement, explain why it is wrong.