2. Oxidation of Food: What a Waste! Fruits and Vegetables oxidised when left in open air Solution: Seal in plastic wrap More radical: Add lemon juice to the cut fruit

3. Oxidation of… Oxidation of nutrients causes increased activity of cells, leading to aging skin Solution: Beauty products? People!

4. What is a redox reaction? Redox – reduction + oxidation Both processes occur simultaneously Hence, one species is oxidised, another is reduced So, what is oxidation, and what is reduction? 3 different versions of the definition:

5. Redox gain of electronsloss of electrons gain in hydrogenloss of hydrogen loss of oxygengain in oxygen ReductionOxidation

6. Gain of oxygen in a species E.g. Mg is oxidized to MgO 2Mg(s) + O 2 (g)  2MgO(s) Loss of hydrogen in a species E.g. HCl is reduced to Cl 2 MnO 2 (s) + 4HCl(aq)  MnCl 2 (aq) + 2H2O(g) + Cl 2 (g)

7. Reduction Loss of oxygen in a species E.g. CuO is reduced to Cu CuO(s) + H 2 (g)  Cu(s) + H 2 O(l) Gain of hydrogen in a species E.g. C 2 H 4 is reduced to C 2 H 6 in the hydrogenation process C 2 H 4 (g) + H 2 (g)  C 2 H 6 (g)

8. Oxidation and Reduction In terms of Electrons (OIL RIG: Oxidation Is Loss, Reduction Is Gain): 2Mg(s) + O 2 (g)  2MgO(s) Oxidation : Mg  Mg 2+ + 2e - (Loss of electrons by Mg) Reduction : O 2 + 4e -  2O 2- (Gain of electrons by O 2 ) The two equations above are known as half- equations.

9. Oxidation numbers For the redox reaction such as S(s) + O 2 (g)  SO 2 (g), would the definition of electron transfer be applicable? A new definition is developed to overcome the problem In terms of Oxidation numbers: Oxidation: Increase in oxidation number in a species Reduction: Decrease in oxidation number in a species The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred.

10. 1.Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H 2, O 2, P 4 = 0 2.In monatomic ions, the oxidation number is equal to the charge on the ion. Li +, Li = +1; Fe 3+, Fe = +3; O 2-, O = -2 3.The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1. Simple rules on oxidation numbers

11. Assign negative charge according to valency AtomIonic chargeOxidation no. FF-F- Always. Most electronegative OO 2- -2Most of the time. Not in cpds with F or peroxides and superoxides ClCl - Not in cpds with O andF. Other oxidation nos include +1, +3, +5, +7 HH+H+ +1Not in metal hydrides, e.g. NaH, where oxidation no is -1 Gp 1Gp 2Gp 3Gp 4Gp 5Gp 6Gp 7 Max oxidation no. +1+2+3+4+5+6+7 Group number as max possible oxidation number

12. Treat a compound as ionic. The most electronegative atom in a molecule is always assigned a negative charge according to its valency to give a full outer shell. E.g. NO 2 +4 + 2 x (-2) = 0

13. 4. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. Deduce the oxidation number of (i)hydrogen in NaH (ii) nitrogen in NH 4 + (iii) chromium in Cr 2 O 7 2- (iv) copper in [CuCl4] 2- (v) cobalt in [Co(H 2 O) 6 ]Cl 2

14. Naming inorganic compounds The oxidation number is inserted immediately after the name of ion. Example Iron(II)chloride, FeCl 2 and iron(III)chloride, FeCl 3 Potassium chromate(VI), K 2 CrO7 and potassium manganate(VII), KMnO4

15. Identify oxidation or reduction reaction Consider the following reactions (1) Cl 2 + 2e  2Cl - (2) Fe 2+  Fe 3+ + e (3) I 2 + 6H2O  2IO 3 - + 12H + + 10e

16. Are these redox reactions? Combination (Synthesis) A + B  AB (a) H 2(g) + O 2(g)  H 2 O (b) CaO (s) + SO 2(g)  CaSO 3(s) Combustion e.g. C x H y + O 2  CO 2 + H 2 O Decomposition ABC  A + B + C 2KClO 3 (s)  2KCl (aq) + 3O 2(aq)

17. Are these redox reactions? Displacement - Single displacement A + BC  AB + CA (a) Cl 2(aq) + NaBr (aq)  NaBr (aq) + Br 2(aq) - Double displacement AB + CD  AD + CB (a) Fe 2 O 3 (s) + HCl (aq)  FeCl 3(aq) + H 2 O (l)

18. Other Non-redox reactions Neutralisation - HA (aq) + BOH (aq )  BA (aq ) + H 2 O (l) - HA (aq) + BO (s)  BA (aq) + H 2 O (l) (a) (b)

19. Precipitation Complex formation Another reaction: Tetraammine copper(II) complex (deep blue solution) ligand

20. Disproportionation A disproportionation reaction is a redox reaction in which one species is simultaneously oxidised and reduced. Example 2H 2 O 2 (aq)  2H 2 O(l) + O 2 (g) Reactant : Oxidation number of O in H 2 O 2 is -1 Product : Oxidation number of O in H 2 O is -2 Oxidation number of O in O 2 is 0 Reduction Oxidation

21. Consider another example: Cl 2 (g) + H 2 O(l)  HOCl(aq) + HCl(aq)

22. Oxidising and Reducing agent An oxidising agent is a substance that brings about oxidation by accepting electrons from the substance it oxidises. It is always reduced in the process. A reducing agent is a substance that brings about reduction by donating electrons to the substance it reduces. It is always oxidised in the process. 2Mg + O 2 → 2MgO Mg is oxidised, and thus is the reducing agent O 2 is reduced, and thus is the oxidising agent

23. List of common Oxidising and Reducing Agents Realise something? H 2 O 2 is both an oxidising and a reducing agent! If a stronger oxidising agent is present, H 2 O 2 is reducing

24. Identifying Redox reactions (a) Consider the reaction S 8 + 12O 2  8SO 3 (i) Does it involve ionic or covalent compounds? (ii) Is this a redox reaction? Explain (iii) What is the oxidising agent and reducing agent?

25. Consider each of the 2 reactions below (a) 2AgNO 3 + Na 2 S  Ag 2 S + 2NaNO 3 (b) 5As 4 O 6 + 8MnO 4 - + 18H 2 O  20 AsO 4 3- + 8 Mn 2+ + 36H + (i) Is this a redox reaction? Explain (ii) What is the oxidising agent and reducing agent?

26. Balancing redox equations The oxidation of Fe 2+ to Fe 3+ by Cr 2 O 7 2- in acid solution. (1) Write the unbalanced equation for the reaction in ionic form. Fe 2+ + Cr 2 O 7 2-  Fe 3+ + Cr 3+ (2) Separate the equation into two half equations Oxidation : Fe 2+  Fe 3+ Reduction: Cr 2 O 7 2- -  Cr 3+ (3) Balance the atoms other than O and H in each half eqn Cr 2 O 7 2-  2Cr 3+ (4) Add H 2 O to balance O atoms and H + to balance H atoms. Cr 2 O 7 2-  2Cr 3+ + 7H 2 O 14H + + Cr 2 O 7 2-  2Cr 3+ + 7H 2 O

27. (5) Add electrons to one side of each half-reaction to balance the charges on the half-reaction. Fe 2+  Fe 3+ + e - 14H + + Cr 2 O 7 2- + 6e -  2Cr 3+ + 7H 2 O (6) If necessary, equalize the number of electrons in the two half- reactions by multiplying the half-reactions by appropriate coefficients. 6Fe 2+  6Fe 3+ + 6e - 14H + + Cr 2 O 7 2- + 6e -  2Cr 3+ + 7H 2 O

28. (7) Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. You should also cancel like species. Oxidation : 6Fe 2+  6Fe 3+ + 6e - Reduction : 14H + + Cr 2 O 7 2- + 6e -  2Cr 3+ + 7H 2 O Overall eqn: 14H + + Cr 2 O 7 2- + 6Fe 2+  6Fe 3+ + 2Cr 3+ + 7H 2 O (8) Verify that the number of atoms and the charges are balanced. 14x1 – 2 + 6x2 = 24 = 6x3 + 2x3 (9) For reactions in basic solutions, add OH - to both sides of the equation for every H + that appears in the final equation. You should combine H + and OH - to make H 2 O.

29. Example Balance the following reaction:

30. Example Write a redox equation for the reduction of acidified manganate(VII) ions and the oxidation of methanol using the balanced half-equations below: (a) 2 H 2 O(l) + CH 3 OH(l)  CO 2 (g) + H 2 O(l) + 6H + (aq) + 6e - (b) MnO 4 - (aq)+ 8H + (aq) + 5e -  Mn 2+ (aq)+ + 4H 2 O(l)

31. Reactivity Series and cell voltage Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Tin Lead Copper Mercury Silver Gold Magnesium Zinc Iron Tin Lead Copper Silver Reactivity decrease Positive electrode Negative electrode Higher Voltage Lower Voltage

32. 2Al + Fe 2 O 3  2Fe + Al 2 O 3 THERMITE REACTION

33. . Uncontrolled spontaneous reaction How can this be harnessed to do work?

34. . Controlled redox reaction Zn  Zn 2+ + 2e − Cu 2+ + 2e −  Cu Zn + Cu 2+  Cu + Zn 2++ overall reaction but electrons are forced outside of the cells

35. Voltaic Cell ΔG = neg Galvanic Battery

36. Voltaic Cells 19.2 spontaneous redox reaction anode oxidation cathode reduction - +

37. Explain what happens in this system. Follow the circulation of one negative charge in the cell.

38. An electric current flows between the anode and cathode because there is a difference in electrical potential between the 2 electrodes, measured by a voltmeter (in volts) : CELL POTENTIAL. The voltaic cell depends on the nature of the electrodes and the ions the concentrations of ions and the temperature at which the celll is operated.

39. CELL POTENTIAL, E Electrons are “driven” from anode to cathode by an electromotive force or emf. Electrons are “driven” from anode to cathode by an electromotive force or emf. For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. Zn and Zn 2+, anode Cu and Cu 2+, cathode 1.10 V 1.0 M

40. CELL POTENTIAL, E For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. is the STANDARD CELL POTENTIAL, E o is the STANDARD CELL POTENTIAL, E o a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C. a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.

41. Cell potentials of selected v cells Metal electrodesCell potenials / V Copper and magnesium2.70 Copper and iron0.78 Lead and zinc0.64 Lead and iron0.32 The further apart the two metals are in the reactivity series, the higher the cell potential.

42. Reactivity Series and cell voltage Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Tin Lead Copper Mercury Silver Gold Magnesium Zinc Iron Tin Lead Copper Silver Reactivity decrease Positive electrode Negative electrode Higher Voltage Lower Voltage

43. Electrolysis Electricity is passed from a battery through a liquid which may be a solution/molten solid. The plates which carry the electricity into the liquid are called electrodes Molten ionic compounds or aqueous solution of ionic compounds that allows electricity to pass through are called electrolytes

44. Electrodes Metal plates or graphite rods that conduct electricity into the electrolyte Eg. Platinum, copper Cathode: Electrode that is connected to the negative terminal of the battery. Postively charged ions, cations, moved towards the cathode Anode: Electrode that is connected to the positive terminal of the battery. Negatively charged ions, anions, moved towards the anode

45. Electrolytic cell

46. Conduction of electricity Conductor is a substance which conducts electricity but is not chemically changed during the conduction Presence of freely moving valence electrons Eg. All metals and graphite Insulator does not allow the passage of electricity. Valence electrons are held in fixed positions Eg. Sulphur, phosphorus, diamond, solid state crystalline salts, wood and glass

47. Electrolytes and non-Electrolytes Electrolytes: Molten ionic compounds or aqueous solution of ionic compounds that allows electricity to pass through and are decomposed in the process Eg. Acids, Alkali, Salts dissolved in water, molten salts Non-electrolytes: Does not allow passage of electricity Eg. Distilled water, alcohol, turpentine, oil, paraffin, organic solvents

48. When electricity is passed through an electrolyte, chemical decomposition occurs This involves the ‘splitting up’ of the electrolyte Since all electrolytes are ionic, composed of positively and negatively charged ions The process: When an electric current pass through the electrolyte, ions in the solution migrate towards the oppositely charged electrode This discharge of ions at the electrodes results in the chemical decomposition of the electrolyte to form its elements.

49. At the anode, negatively charged ions lose their electron(s) to the anode (connected to positive terminal of battery) to form neutral atoms. The negatively charged ions are said to be oxidised and discharged at the anode. Oxidation occured at the anode. At the cathode, positively charged ions gain electron(s) from the cathode (connected to negative terminal of battery) to form neutral atoms. The positively charged ions are said to be reduced and discharged at the cathode. Reduction occured at the cathode.

50. Rules for Predicting Selective Discharge of Cations

51. Rules for Predicting Selective Discharge of Anions

52. Electrolysis of Concentrated Hydrochloric Acid Carbon rods as electrodes

53. Electrolysis of Molten Compounds Many ionic compounds are binary compounds. A binary compound is a compound containing only 2 elements. It contains a metal cation and a non-metal anion. The electrolysis of a molten binary compound will yield a metal and a non-metal as products.

54. Electrolysis of Molten Lead(II) Bromide Carbon rods as electrodes

55. Electrolysis of Molten Sodium Chloride

56. Electrolysis of Water 19.8

57. Chemistry In Action: Dental Filling Discomfort Hg 2 /Ag 2 Hg 3 0.85 V 2+ Sn /Ag 3 Sn -0.05 V 2+ Sn /Ag 3 Sn -0.05 V 2+

58. Corrosion

59. Cathodic Protection of an Iron Storage Tank

60. Simple Voltaic CellsElectrolytic cell Chemical energy changed to electrical energy Electrical energy changed to chemical energy 2 electrodes, 2 different metals 2 electrodes, same or different metals

61. Identifying redox reactions Consider the following reactions (i) 2FeCl 2 (s) + Cl 2 (g)  2FeCl 3 (s) (ii) Mn(NO 3 ) 2 (s)  MnO 2 (s) + 2NO 2 (g) (iii) (NH 4 ) 2 Cr 2 O 7 (s)  Cr 2 O 3 (s) + 4H 2 O(g) + N 2 (g) What if there’s no change in the oxidation numbers during the chemical reaction?