2. Unit 12: Redox and Electrochemistry - RB Topic 11

3. *What kind of chemical reaction is responsible for burning, rusting, bleaching clothes, powering your cell phone as well as your body, and giving fireflies their glow? Oxidation-Reduction Reactions, also called REDOX

4. I. What are Redox Reactions? involve the transfer (gain/loss) of ELECTRONS (e – ) between different atoms

5. OXidationREDuction losing electrons charge goes up Ex: Na + Cl  Na + Cl gaining electrons charge is reduced Ex:

6. *CONSERVATION OF CHARGE* oxidation and reduction happen simultaneously (at the same time) – you can’t have one without the other  charges on both sides of arrow MUST BE EQUAL

7. Remember: LEO the lion says GER LOSING ELECTRONS is OXIDATION GAINING ELECTRONS is REDUCTION

8. II. Determining if a reaction is a redox reaction If the oxidation number (charge) of an atom CHANGES from reactant to product, a redox reaction has occurred

9. *Note: Double replacement reactions are NOT redox reactions Single replacement reactions are ALWAYS redox reactions Ex: CaCl 2 + Na 2 S  CaS + 2NaCl Ca + MgS  Mg + CaS

10. A. Assigning Oxidation Numbers ( listed on the Periodic Table) 1. If a species is all by itself (not combined with a different element), its oxidation number is ZERO Ex: Mg O 2 S 2 Cu H 2

11. 2. In a polyatomic ion, the oxidation #’s of the elements add up to equal the CHARGE on the ion *first element is always + *last element is always - NH 4 + ClO –

12. SO 3 2– SO 4 2–

13. 3. In a compound, the oxidation numbers of the elements add up to ZERO (compounds are neutral) *first element is always + *last element is always – NaClNH 4 Cl

14. H 2 SO 3

15. More Examples: a.) What is the oxidation number of chlorine in iron(III) chloride? b.)What is the oxidation state of chromium in potassium chromate?

16. Skill 1

17. B. Checking for Changes in Oxidation Numbers – if they change…it’s REDOX 1. start by assigning ox. #s 2 Na (s) + Cl 2 (g) → 2NaCl (s) The oxidation # on Na changes from 0 to +1 Na loses electrons Na is oxidized Na is the reducing agent

18. 2 Na (s) + Cl 2 (g) → 2NaCl (s) The oxidation # on Cl changes from 0 to –1 Cl 2 gains electrons Cl 2 is reduced Cl 2 is the oxidizing agent

19. 2Mg (s) + O 2 (g) → 2MgO (s) The oxidation # on Mg changes from 0 to +2 Mg loses electrons Mg is oxidized Mg is the reducing agent The oxidation # on O changes from 0 to –2 O 2 gains electrons O 2 is reduced O 2 is the oxidizing agent

20. Skill 2

21. Ticket Given the following: 4 Mn + 7 O 2  2 Mn 2 O 7 (a)Assign oxidation numbers to each symbol in the reaction (b)Write the formula of the substance that is oxidized (c)Write the formula of the substance that is reduced

22. III.Writing Half-Reactions * Remember: conservation of MASS and CHARGE 1. Write oxidation #s 2. Write in electrons gained (reactant) or lost (product) and include COEFFICIENTS 3. Make sure # of atoms and total charge STAYS THE SAME (conservation of MASS and CHARGE) 2 Na + Cl 2 → 2 NaCl ___ Na + ___Cl 2 → ___ NaCl Oxidation half-reaction: shows the species being oxidized and the electrons that are lost (PRODUCED) Ex: Mg 2+ + 2e – → Reduction half-reaction: shows the electrons that are gained and the species being reduced Ex: 2I –  I 2 + 2e –

23. Reduction half-reaction: shows the electrons that are gained and the element being reduced Ex: Mg 2+ + 2e – → Oxidation half-reaction: shows the element being oxidized and the electrons that are lost Ex: 2I –  I 2 + 2e – III.Writing Half-Reactions 1. Write oxidation #s 2. Write in electrons gained (reactant) or lost (product) and include COEFFICIENTS 3. Make sure # of atoms and total charge STAYS THE SAME (conservation of MASS and CHARGE) 2 Na + Cl 2 → 2 NaCl

24. Sample Problems: Write reduction and oxidation half- reactions for each of the following: (a) 2Fe 3+ + Ni  2Fe 2+ + Ni 2+

25. (b) Br 2 + Hg  Hg 2+ + 2Br –

26. (c) Cu + 2Ag +  2Ag + Cu 2+

27. (d) Sn 4+ + H 2  Sn + H +

28. Skill 3

29. In the presence of oxygen, iron forms iron (III) oxide, more commonly known as rust, through a reduction-oxidation reaction 4Fe + 3O 2  2Fe 2 O 3

31. What causes fruit to turn brown? Oxidation of compounds in fruit after exposure to oxygen in the air causes a change of color (browning) of fruit

32. Oxidation (losing electrons) is a process that can be very damaging to living things. results in the creation of free radicals, which are chemical species that have unpaired electrons (they are unhappy – they don’t have a completed valence shell) free radicals will rip electrons away from other compounds in living things, such as proteins or molecules of DNA, causing damage which can lead to cancer Antioxidants get rid of free radicals – Vitamins A, C, and E are antioxidants

33. IV. Electrochemical Cells A. Use REDOX reactions to convert chemical energy into electrical energy OR convert electrical energy into chemical energy

34. B. Two types of Electrochemical Cells 1. Voltaic /Galvanic (battery) convert chemical  electrical energy Use a chemical reaction to produce RELEASE electricity (moving electrons) 2. Electrolytic Cells convert electricity  chemical energy use electricity to make a chemical reaction to occur)

35. 1. Voltaic (Galvanic) Cells = BATTERIES a. SPONTANEOUS redox reactions – NO energy put in (energy is released) ***Use Table J to determine what is oxidized and what is reduced Higher on Table J is OXIDIZED (LEO) Lower on Table J is REDUCED (GER)

36. b. Diagram of a voltaic cell using Zn and Zn(NO 3 ) 2 with Cu and Cu(NO 3 ) 2

37. *Parts to Know*  Electrodes: the places where oxidation or reduction happens o Anode: the site of oxidation (An Ox) o Cathode: the site of reduction (Red Cat)  Wire: electrons flow through the wire from the anode (–) to the cathode (+)  Salt bridge: allows IONS to flow between the electrodes to keep the charges balanced

38. The electrodes must be separated in order to produce an electric current (flow of electrons). The energy present in the flowing electrons (ELECTRICITY) is captured and used to power other processes.

39. c. Voltaic cell problems: 1.Look on Table J and find which element is higher – this element is OXIDIZED (On Table J – electrons flow DOWNHILL, spontaneously) 2.Under each beaker write “oxidation” or “reduction” 3.Label the oxidation electrode “ANODE” (An Ox) 4.Label the reduction electrode “CATHODE” (Red Cat)

40. 5. Place a (–) charge on the anode 6. Place a (+) charge on the cathode 7.Draw in the direction of electron flow (from ANODE (–) to CATHODE (+)) 8.Write the half reactions under the correct beakers

41. 9. The salt bridge allows for ions to flow between the electrodes Positive (+) ions flow toward the CATHODE (to balance the negative electrons) Negative (–) ions flow toward the ANODE (to replace the electrons that are leaving)

42. b. Diagram of a voltaic cell using Zn and Zn(NO 3 ) 2 with Cu and Cu(NO 3 ) 2

43. Label the parts of this voltaic cell

44. How many of you have had cavities filled? …if you do, you have the potential to make a tiny battery IN YOUR MOUTH! Cavities are filled with a mixture of metals including zinc (Zn), tin (Sn), copper (Cu), and silver (Ag) If you bite down on a piece of aluminum foil, the saliva in your mouth, the aluminum foil, and the filling make a little voltaic cell that produces a tiny current that travels through your tooth to the nerve below the filling…it’s a bit UNPLEASANT!

45. Electrochemical Cells Lab Lemon battery lab

46. 2. Electrolytic Cells a. NONSPONTANEOUS redox reactions – need to put in energy b. The species more likely to lose electrons is forced to gain electrons (On Table J – electrons are moving UPHILL, nonspontaneously) c. Electrons still flow from the anode to the cathode, but now the signs are reversed The anode is (+) and the cathode is (-) (electrons are being forced to travel to the negative electrode)

47. d. Used for 1. Electrolysis: using electricity to break apart (lyse) compounds into their elements Ex: Obtaining active elements from compounds (a) complete and balance following electrolysis reaction. (b) assign oxidation numbers to each element ___NaCl ( l )

48. ____H 2 O ( l )

50. 2. Electrolytic cells can also be used for electroplating – putting a metal coating on something

52. Comparing & Contrasting Voltaic and Electrolytic Cells Voltaic Cells (Batteries)Electrolytic Cells SPONTANEOUS redox rxn energy is RELEASED! (RELEASES electricity) nonspontaneous redox rxn energy is needed (uses electricity)

53. Comparing & Contrasting Voltaic and Electrolytic Cells Voltaic Cells (Batteries)Electrolytic Cells) redox reactions anode is where oxidation happens (An Ox) anode loses mass cathode is where reduction happens (Red Cat) cathode gains mass electrons flow through wire from anode to cathode

54. p. 186 #s 37-41, 51-52, 65-69, 71-74, 76-86, 88-91 37. 3 38. 2 39. 3 40. 2 41. 4 51. 2 52. 2 65. 1 66. 4 67. 2 68. 4 69. 3 76. 2 77. 3 78. 2 79. 2 80. 3 81. 2 71. 1 72. 1 73. 3 74. 4 82. 4 83. 1 84. 4 85. 2 86. 1 91. The anode loses mass, the cathode gains mass