Lesson 1- Introduction to Redox Reactions

1. Unit 7- Redox Chemistry and Electrochemistry Lesson 1- Introduction to Redox Reactions • Define oxidation and reduction in terms of oxygen loss/gain • Define redox in terms of electron transfer. • Construct redox half equations • Assign Oxidation States

2. What is a redox reaction? oxygen removed reduction lead oxide carbon monoxide +  + carbon lead oxygen added oxidation redox = reduction and oxidation Oxidation is the addition of oxygen to a substance and reduction is the removal of oxygen from a substance. Which substances are oxidized and reduced in this reaction? Reduction and oxidation always take place together. Why is this type of reaction called a redox reaction?

3. Redox reactants – oxidized or reduced?

4. Redox and electrons  magnesium + oxygen magnesium oxide +  2Mg(s) O2(g) 2MgO(s) Magnesium burns in oxygen to form magnesium oxide. It is obvious that the magnesium has been oxidized, but what has happened to the oxygen? A redox reaction can also be explained in terms of the gain or loss of electrons. What happens to the atoms and electrons in this reaction?

5. Oxidation and electron loss oxidized (electrons lost) +  Mg O Mg2+ O2- Oxidation is the loss of electrons. When magnesium burns in oxygen to form magnesium oxide, what happens to magnesium and its electrons? • The magnesium has been oxidized. • The Mg atom has lost 2 electrons to form a Mg2+ion.

6. Reduction and electron gain reduced (electrons gained) +  Mg O Mg2+ O2- Reduction is the gain of electrons. When magnesium burns in oxygen to form magnesium oxide, what happens to oxygen and its electrons? • The oxygen has been reduced. • The O atom has gained 2 electrons to form a O2-ion.

7. Redox and OILRIG An easy way to remember what happens to the electrons during oxidation and reduction is to think… OILRIG!

8. Using OILRIG Oxidation Is Loss of electrons Reduction Is Gain of electrons What does OILRIG stand for in terms of redox reactions?

9. What is a half-equation?  magnesium + oxygen magnesium oxide +  2Mg(s) O2 (g) 2MgO (s) oxidation: Mg  Mg2+ + 2e- reduction: O2 + 4e-  2O2- Redox reactions involve the transfer of electrons. Equations written to show what happens to the electrons during oxidation and reduction are called half-equations. What are the half-equations for the oxidation and reduction processes in this reaction?

10. What does each half-equation show?

11. Redox reactions – summary

12. Questions • Explain why the reaction between calcium and oxygen to form calcium oxide is an example of a redox reaction. • Define oxidation and reduction. In terms of oxygen and electrons. • Balance the following half equations Br2(l) + _e¯ → 2Br¯(aq) Ag+(aq) + _e¯ → Ag(s) Fe3+(aq) + _e¯ → Fe2+(aq) I2(s) + _e¯ → 2I¯(aq) Cu +(aq) + _e¯ → Cu(s) Cu2+(aq) + _e¯ → Cu(s)

13. Oxidation numbers • We can use oxidation numbers to decide whether a something has been oxidised or reduced. • Oxidation numbers show how many electrons are gained or lost by an element when atoms turn into ions. • The oxidation numbers of the elements are zero. Oxidation numbers also distinguish between the compounds of elements such as iron which can exist in more than one oxidation state. In iron (II) chloride, the Roman number II shows that the iron is in oxidation state +2.

14. Oxidation State/Number Rules: • The oxidation number of an element is zero. • The oxidation number of an ion is equal to the charge of the ion. • Hydrogen has an oxidation number of +1 (this doesn't apply to hydrides when hydrogen is -1). • Oxygen has an oxidation number of -2 (except in peroxides when it is -1). • Fluorine always has an oxidation number of -1 • It is then simply a matter of adding together all of the oxidation numbers of the elements in a compound and making sure that the total is = 0 or equal to the charge on a compound ion.

15. Assigning Oxidation State- Examples Cr2O3 KMnO4 HNO3

16. Assign oxidation numbers to each element in each of the following compounds: 1. KMnO4 2. NiO2 3. P4O6 4. Fe3O4 5. SF4 6. XeOF4 7. CO 8. Na2C2O4 9. As2O3 10. NaBiO3 11. Mg2P2O7 12. Hg2Cl2

17. Oxidation Numbers- Charged Species Cr2O72- KMnO4 -

18. Identifying Redox Reactions Example 1: Balanced Equations: Zn + CuCl2 → ZnCl2 + Cu balanced equation Zn + Cu2+→ Zn2+ + Cu ionic equation Note: chloride is a spectator ion Half Equations: Zn → Zn2+ +2 e- oxidation Cu2++ 2e-→ Cu reduction

19. Example 2: Identify redox processes AgNO3 + NaCl → AgCl + NaNO3

20. Use oxidation numbers to identify whether the following processes, represented by equations, involve oxidation (O), reduction (R), both or neither. Give reasons for your decisions. (a) F2 +2e → 2F- (b) MnO2 + 2H2O → MnO4- + 4H+ + 3e (c) H2C2O4 → 2CO2 +2H+ + 2e (d) 2CrO42- + 2H+→ Cr2O72- + H2O (e) C2H4 + Cl2→ C2H4Cl2 (f) Zn → Zn2+ + 2e (g) Fe2+ + 2e → Fe

21. Example 3: N2 + 3H2 → 2NH3 Oxidation is a increase in Oxidation Number Reduction is a decrease in Oxidation Number

22. Task Complete Worksheet 1:

24. Unit 7- Redox Chemistry and Electrochemistry Lesson 2- Electrolysis of molten ionic compounds • Properties of ionic compounds • Define the term electrolysis • Describe the electrode products in the electrolysis of -Molten lead(II) bromide

25. Starter • Say whether the following three reactions are redox reactions or not, by showing the oxidation states of each of the elements involved: a) Zn + 2HCl → ZnCl2 + H2 b) CuO + 2HCl → CuCl2 + H2O c) MnO2 + 4 HCl → MnCl2 + Cl2 + 2 H2O d) Cl2 + H2O → HCl + HOCl

26. What are ionic compounds? Ionic compounds are made up of positive metal ions and negative non-metal ions. What ions are in sodium chloride? positive sodium ions negative chloride ions The positive and negative ions in an ionic compound attract each other strongly. It takes a lot of energy to separate them. How does structure affect the properties of ionic compounds?

28. Do dissolved ionic compounds conduct?

29. Properties of ionic compounds

30. Does it conduct electricity?

31. Electrolysis A process in which a chemical reaction is caused by the passage of an electric current. A molten ionic compound can be split in to its elements. eg PbCl2 into Lead metal and Chlorine gas.

32. Electrode- The point at where the electric current enters and leaves a current a battery. Electrolyte- an ionic compound which will conduct electricity when it is molten or dissolved in water; electrolytes will not conduct electricity when solid.

33. What is electrolysis? An ionic compound conducts electricity when it is molten or in solution. The current causes the ionic compound to split up and form new substances. This process is called electrolysis, and means “splitting by electricity”. Electrolysis has many uses, including: • purifying copper • plating metals with silver and gold • extracting reactive metals, such as aluminium • making chlorine, hydrogen and sodium hydroxide.

34. What happens during electrolysis? heat In electrolysis, the substance that the current passes through and splits up is called the electrolyte. The electrolyte contains positive and negative ions. What happens to these ions during electrolysis? Negative ions move to the positive electrode (Anode) and lose electrons. This is oxidation. Positive ions move to the negative electrode (Cathode) and gain electrons. This isreduction.

35. Electrolysis of molten lead bromide

36. Electrolysis of molten PbBr2 – redox equations leadbromidelead + bromine PbBr2(l)Pb(l) + Br2(g) What redox processes occur at the electrodes during the electrolysis of molten lead bromide (PbBr2)? At the negative electrode: Pb2+ + 2e- Pb (reduction) At the positive electrode: 2Br-  Br2 + 2e- (oxidation) What is the overall equation for the electrolysis of molten lead bromide?

37. Electrolysis of molten PbBr2 – summary

38. Electrolysis problems 1 1) Define electrolysis. (1) 2) a) Explain monatomic and simple molecular substances do not conduct electricity. (2) b) Explain why metals conduct electricity. (3) c) Explain why ionic substances do not conduct electricity as solids but do when molten or dissolved. (3) 3) When each of the following ionic substances is melted and electrolysed, layout your answers clearly: • predict what would be produced at the positive electrode • give the half equation for the reaction at the positive electrode • predict what would be produced at the negative electrode • give the half equation for the reaction at the negative electrode a) molten lead bromide

39. Electrolysis problems 1 • Define electrolysis. A)The use of an electrical current to split up the components of a molten ionic compound or the components within a solution (1) • a) Explain monatomic and simple molecular substances do not conduct electricity. A)Molecular substances do not consist of ions or electrons that are able to move freely and conduct electricity. (2) • Explain why metals conduct electricity. (3) A)Metals consist of free electrons which are able to move and transport charge. • Explain why ionic substances do not conduct electricity as solids but do when molten or dissolved. • The ions within a solid ionic substance are unable to move and transport electrical current, when molten or dissolved the ions are able to move and transport electrical charge

40. 3) When each of the following ionic substances is melted and electrolysed, layout your answers clearly: • predict what would be produced at the positive electrode • give the half equation for the reaction at the positive electrode • predict what would be produced at the negative electrode • give the half equation for the reaction at the negative electrode • molten lead bromide At the negative electrode: Product: Lead Half Equation: Pb2+ + 2e-  Pb (reduction) At the positive electrode: Product: Bromine Gas Half Equation: 2Br-  Br2 + 2e- (oxidation)

41. Unit 7- Redox Chemistry and Electrochemistry Lesson 3- Electrolysis of Ions in Solution • Define the term electrolysis • Describe the electrode products in the electrolysis of -Molten lead(II) bromide -Concentrated hydrochloric acid -Concentrated aqueous sodium chloride. -Between inert electrodes (platinum or carbon).

42. Starter Draw a rough sketch of the equipment used during electrolysis. Include: • Electrolyte • Electrodes • Anode • Cathode

44. The electrolysis of HCl • Task • Can you construct balanced half equations for this process. • Combine the half equation to give the full equation of what is happening. • What are the products of this reaction.

45. At the negative electrode: Product: Hydrogen Half Equation: 2H+ + 2e- H2(reduction) At the positive electrode: Product: Chlorine Half Equation: 2Cl- -> Cl2 + 2e- (Oxidation) Overall: 2HCl -> H2 + Cl2

46. Questions • Electrolysis: Splitting up with electricity. • This requires a liquid, or ________, which will conduct electricity. • Electrolytes are usually free _____ dissolved in water. • Electrolytes can also be ________ ionic substances. • The _________ is an electron pump, taking electrons away from the positive ______ to the negative _______. • 5. Ions gain or lose electrons at the ______ and _____ atoms or molecules are formed. • 6. The are ______ often made of _____ materials such as platinum and carbon neutral electrolyte electricity cathode inert electrodes electrodes anode ions molten

47. The electrolysis of Concentrated aqueous sodium chloride. What would you expect the products of this reaction would be?

48. Electrolysis of NaCl solution

49. Products of electrolysis of NaCl solution The electrolysis of sodium chloride solution produces three very useful products: • Chlorine used for killing bacteria in water, for bleach and making plastics like PVC. • Hydrogen used for making margarine and fertilizers, and for rocket fuel. • Sodium hydroxide used in many chemical reactions, such as making soap,neutralizing acids and making paper. Chlorine is expected as a product of this process but hydrogen and sodium hydroxide are surprising products. What happens at the electrodes to form these products?

50. How does the chlorine form? In the electrolysis of NaCl solution, the negative chloride ions (Cl-) are attracted to the positive electrode. Here, the Cl- ionslose electrons to make chlorine atoms, which then form chlorine molecules (Cl2). oxidized • Are the Cl- ions oxidized or reduced? • How many electrons are lost by each Cl- ion? • How many Cl- ions join to make a Cl2 molecule? • What is the half-equation for this redox process? one two 2Cl-  Cl2 + 2e-(oxidation)