1. Chapter 20 – Redox Reactions One of the earliest recognized chemical reactions were with oxygen. Some substances would combine with oxygen, and some would release oxygen. If a substance would gain oxygen, it was said that the substance was oxidized. If the substance lost oxygen, it was said that it was reduced. Since then, reduction-oxidation reactions (redox reactions) have been defined to explain why these substances lost or gained oxygen.

2. Redox Reactions Oxidation (Old School Definition); Any substance that gained oxygen in a chemical reaction. 4Fe (s) + 3O 2(g)  2Fe 2 O 3(s) We can say that the iron was oxidized.

3. Redox Reactions Reduction (Old School Definition); Any substance that loses oxygen in a chemical reaction. 2Fe 2 O 3(s) +3C (s)  4Fe (s) +3 CO 2(g) We can say that the iron (III) oxidize was reduced.

4. Redox Reactions Oxidation (New School Definition); Any substance that loses electrons in a chemical reaction. 3Mg (s) + N 2(g)  Mg 3 N 2(s) In this chemical reaction, Mg goes from zero charge as a reactant to a +2 charge as a product. Mg (s)  Mg +2 (aq) + 2e -

5. Redox Reactions Reduction (New School Definition); Any substance that gains electrons in a chemical reaction. 3Mg (s) + N 2(g)  Mg 3 N 2(s) In this chemical reaction, N goes from zero charge as a reactant to a -3 charge as a product. N 2(s) + 6e -  2N -3 (aq)

6. Redox Reactions Reduction-Oxidation (Redox) Reactions Redox reactions occur when one atom loses electrons, while another gains electrons. We can tell if electrons have been transferred by looking at the individual charge (oxidation number), of each atom.

7. Redox Reactions Rules for Assigning Oxidation Numbers 1.Monoatomic Ions have the same sign and charge of its ionic charge.

8. Redox Reactions Rules for Assigning Oxidation Numbers 2.The oxidation number for hydrogen in a compound is +1 except when it is part of a metal hydride.

9. Redox Reactions Rules for Assigning Oxidation Numbers 3.The oxidation number of oxygen in a compound is -2 unless it is part of a peroxide where it is -1. Oxygen can have a -1 charge if it is part of a compound where there is an element more electronegative than it.

10. Redox Reactions Rules for Assigning Oxidation Numbers 4.The oxidation of atoms in their elemental forms is zero.

11. Redox Reactions Rules for Assigning Oxidation Numbers 5.For neutral compounds, the sum of the oxidation numbers must equal zero.

12. Redox Reactions Rules for Assigning Oxidation Numbers 6.For polyatomic ions, the sum of the oxidation numbers must equal the charge of the polyatomic ion.

13. Redox Reactions Assigning Oxidation Numbers Assign an oxidation number for each element in the following compounds; 1.) SO 2

14. Redox Reactions Assigning Oxidation Numbers Assign an oxidation number for each element in the following compounds; 2.) CO 3 -2

15. Redox Reactions Assigning Oxidation Numbers Assign an oxidation number for each element in the following compounds; 3.) Na 2 SO 4

16. Redox Reactions Assigning Oxidation Numbers Assign an oxidation number for each element in the following compounds; 4.) (NH 4 ) 2 S

17. Redox Reactions Assigning Oxidation Numbers Assign an oxidation number for each element in the following compounds; 5.) Fe

18. Redox Reactions Rules for Balancing Redox Reactions 1.) Divide the chemical equation into 2 half-reactions, one for the oxidation and one for the reduction.

19. Redox Reactions Rules for Balancing Redox Reactions 2.) Balance each half reaction; a.) First, balance the elements other than H and O. b.) Next, balance the O atoms by adding H 2 O as needed. c.) Then, balance the H atoms by adding H + as needed. d.) Finally, balance the charge by adding e - ’s as needed.

20. Redox Reactions Rules for Balancing Redox Reactions 3.) Multiply the half-reactions by integers so that the number of e - ’s lost in one half-reaction equals the number of e - ’s gained in the other.

21. Redox Reactions Rules for Balancing Redox Reactions 4.) Add the two half-reactions and combine and cancel out molecules.

22. Redox Reactions For reactions that occur in basic solution – o Balance as before but add an equal amount of OH - ions to each side of the half reaction to cancel out the H +. o Complete and balance the follow redox reaction that occurs in a basic solution; CN - (aq) + MnO 4(aq)  CNO - (aq) + MnO 2(s)

23. Redox Reactions CN - (aq) + MnO 4(aq)  CNO - (aq) + MnO 2(s)

24. Voltaic Cells (Galvanic Cells) A device in which the transfer of electrons takes place through an external path. Oxidation occurs at the anode. Reduction occurs at the cathode.

25. Voltaic Cells (Galvanic Cells) A device in which the transfer of electrons takes place through an external path. The Zn strip serves as the anode. The Cu strip serves as the cathode.

26. Voltaic Cells (Galvanic Cells) A device in which the transfer of electrons takes place through an external path. The salt bridge serves as a way to alleviate the build-up of electric charge. Why?

27. Voltaic Cells (Galvanic Cells) The two half-reactions in a voltaic cell are Zn (s)  Zn 2+ (aq) + 2e - ClO 3 - (aq) + 6H + (aq)  Cl - (aq) + 3H 2 O (l) a.) Indicate which reaction occurs at the anode and which occurs at the cathode.

28. Voltaic Cells (Galvanic Cells) The two half-reactions in a voltaic cell are Zn (s)  Zn 2+ (aq) + 2e - ClO 3 - (aq) + 6H + (aq)  Cl - (aq) + 3H 2 O (l) b.) Which electrode is consumed in the reaction? c.) Which electrode is positive?

29. Voltaic Cells (Galvanic Cells) How do we measure the potential energy difference between the anode and the the cathode? o The electromotive force, emf, is the potential difference between the anode and the cathode. 1 Volt = 1 Joule Coulomb

30. Voltaic Cells (Galvanic Cells) Electromotive Force o emf = E cell = cell potential o Under standard conditions – (25°C, 1 atm, and 1M) standard emf = E° cell = standard cell potential 1 Volt = 1 Joule Coulomb

31. Voltaic Cells (Galvanic Cells) Electromotive Force o E° cell = E° red (cathode) – E° red (anode) * (Appendix E) o Calculate the standard cell potential for a Zn / Cu +2 voltaic cell?

32. Voltaic Cells (Galvanic Cells) Electromotive Force o E° cell = E° red (cathode) – E° red (anode) * (Appendix E) o Using Appendix E, calculate the standard emf for the following voltaic cell; Cr 2 O 7 2- (aq) + 14H + (aq) + 6I - (aq)  2Cr 3+ (aq) + 3I 2(s) + 7H 2 O (l)

33. Voltaic Cells (Galvanic Cells) Electromotive Force o E° cell = E° red (cathode) – E° red (anode) * (Appendix E) o A voltaic cell is based on the following two half- reactions; Cd 2+ (aq) + 2e -  Cd (s) Sn 2+ (aq) + 2e -  Sn (s) a.) Use appendix E to determine the half-reactions that occur at the anode and the cathode.

34. Voltaic Cells (Galvanic Cells) Electromotive Force o E° cell = E° red (cathode) – E° red (anode) * (Appendix E) o A voltaic cell is based on the following two half- reactions; Cd 2+ (aq) + 2e -  Cd (s) Sn 2+ (aq) + 2e -  Sn (s) b.) Calculate the standard cell potential of this voltaic cell.

35. Free Energy (ΔG) and Redox Reactions ΔG = -nFE n = number of electrons transferred in the reaction. F = Faraday’s Constant (96,485 J/V.mol) E = emf (cell potential) If ΔG has a negative sign, then what can we conclude about the spontaneity of the chemical reaction?