1. Oxidation and Reduction Reactions Chapters 20 and 21

2. The electrical charge that an atom or ion has- or appears to have- when combined with other elements in a compound or polyatomic ion

3. Free elements= 0 charge Group 1 metals= +1 Li +, Na +, K + Group 2 metals= +2 Ca 2+, Mg 2+ Halogens= -1 F -, Cl -

4. Oxygen: always -2 except peroxides (such as H 2 O 2 ) where each O is -1 OR in combination with fluorine, OF 2, where O is +2 Hydrogen: always +1 except metal hydrides (where it is -1) LiH, CaH 2

5. The total of the oxidation numbers in neutral compounds must equal 0 The oxidation numbers of all the atoms in an ion, (including polyatomics) must add up to the charge on the ion

7. REDUCTION: gain electrons GER Oxidation number is reduced OXIDATION: lose electrons LEO Oxidation number gets bigger

9. 2Mg + O 2  2Mg +2 + 2O -2 Oxidation (lose e - ) Reduction (gain e - )

11. An element loses electrons becomes oxidized Charge becomes MORE POSITIVE This particle reduces another by letting it “take” its electrons reducing agent Group 1: STRONG reducing agents strong tendency to lose electrons

12. L ose R aise R educing E lectrons O xidation A gent O xidation N umber

13. An element gains electrons Becomes reduced- lower oxidation # Oxidation number decreases Becomes more negative This particle oxidizes another particle by removing an e - from it Oxidizing agent Group 17: STRONG OXIDIZING AGENTS Strong tendency to accept e - and become reduced

14. G ain L ower O xidizing E lectrons O xidation A gent R eduction N umber

16. 1.Assign oxidation numbers to all elements 2.Separate the rxn into half-rxns 3.Equalize the number of elements  Conservation of mass!!

17. 4. A) equalize the charge by adding e - in LEO. Do the same for GER B) balance the number of e - between LEO and GER by multiplying by common # 5. Combine 1/2 rxns into skeleton equation 6. Balance the rest of the main equation conservation of mass

18. Single replacement, decomposition, synthesis: ALWAYS redox Double replacement: NOT redox

19. For any TWO METALS in an activity series (TABLE J), the more active metal is the more readily oxidized Metals want to lose electrons Halogens: the more active is more easily reduced due to higher electronegativity Nonmetals want to gain electrons

20. Beaker 1 BeforeCu +2 solution Zn o strip Beaker 1 After Cu o is deposited on Zn o strip Cu +2 solution (blue) turns clear

21. Atom (Zn o ) comes before the ion (Cu +2 ) on Table J  SPONTANEOUS Metals WANT to lose electrons, so more active metal is oxidized Ion is “forced” to gain electrons and become reduced

22. Used to determine the direction of a spontaneous reaction Recall: if one reactant is reduced, the other MUST BE OXIDIZED

23. Strong tendency to lose electrons Undergo oxidation (LEO) Arranged from best oxidizers at the top Element higher up starts as the atom (0 charge) and gets oxidized (becomes + ion)

24. Strong tendency to gain electrons (undergo reduction  become - ) Arranged from best reducers at the top

25. The study of the conversion of chemical energy to electrical energy Electrochemical cell- converts chemical energy into electrical energy or electrical energy into chemical energy

26. LEOxidation --------> GEReduction IDEA: put the flow of electrons through a wire! Electricity!!!

27. cathode anode  Electron flow Salt bridge An Ox: Oxidation occurs at the anode Red Cat: Reduction occurs at the cathode Half cell

28. Salt bridge Zn 2+ Cu 2+ Na + Zn Cu SO 4 2– Voltmeter Anode (–) (+) Slide from University of Washington Chem Dept.

29. Oxidation half-reaction Zn(s) Salt bridge Zn 2+ Cu 2+ Na + Zn Cu SO 4 2– Zn 2+ (aq) + 2e – Voltmeter e–e– Anode (–) (+) Slide from University of Washington Chem Dept.

30. Zn 2+ Zn Oxidation half-reaction Zn(s) Salt bridge Zn 2+ Cu 2+ Na + Zn Cu SO 4 2– Zn 2+ (aq) + 2e – Voltmeter e–e– 2e – lost per Zn atom oxidized Anode (–) (+) e–e– Slide from University of Washington Chem Dept.

31. Zn 2+ Zn Oxidation half-reaction Reduction half-reaction Cu 2+ (aq) + 2e – Zn(s) Salt bridge Zn 2+ Cu 2+ Na + Zn Cu SO 4 2– Zn 2+ (aq) + 2e – Cu(s) Voltmeter e–e– e–e– 2e – lost per Zn atom oxidized Anode (–) Cathode (+) e–e– Slide from University of Washington Chem Dept.

32. Cu 2+ e–e– Cu 2e – gained per Cu 2+ ion reduced Zn 2+ Zn Oxidation half-reaction Reduction half-reaction Cu 2+ (aq) + 2e – Zn(s) Salt bridge Anode (–) Cathode (+) Zn 2+ Cu 2+ Na + Zn Cu SO 4 2– Zn 2+ (aq) + 2e – Cu(s) Voltmeter e–e– e–e– 2e – lost per Zn atom oxidized e–e– Slide from University of Washington Chem Dept.

33. Cu 2+ e–e– Cu 2e – gained per Cu 2+ ion reduced Zn 2+ Zn Oxidation half-reaction Reduction half-reaction Overall (cell) reaction Zn(s) + Cu 2+ (aq) Cu 2+ (aq) + 2e – Zn(s) Salt bridge Zn 2+ Cu 2+ Na + Zn Cu SO 4 2– Zn 2+ (aq) + 2e – Cu(s) Zn 2+ (aq) + Cu(s) Voltmeter e–e– e–e– Anode (–) Cathode (+) 2e – lost per Zn atom oxidized e–e– Slide from University of Washington Chem Dept.

34. Salt Bridge: when e - move one solution will become very negative, e - won’t want to go there. The salt bridge allows for ION MIGRATION Without a salt bridge, the circuit is incomplete and e - cannot flow through the wire.

35. The process by which electrical energy is used to make nonspontaneous redox rxns proceed

36. Anode  oxidation (+) Cathode  reduction (-) Electrons flow from LEO  GER (+) (-) **NONSPONTANEOUS Must supply electricity to force the rxn to occur  electrons are not attracted to negatively charged substances

37. Spontaneous redox rxns (salt bridge) converts chemical to electrical energy e - flows from anode to cathode (-) (+) ox red Electric current used to drive a nonspontaneous redox reaction Requires electrical energy to produce a chemical change ELECTROLYSIS/electroplating e - flow from the anode to cathode (+) (-) Voltaic CellElectrolytic Cell

38. 2 H 2 O + electricity  2H 2 + O 2 Main clue that you have electrolytic cell

39. 2 KCl + electricity  2 K o + Cl 2 o Cathode (-): 2K + + 2e -  2K Anode (+): 2Cl -  Cl 2 + 2e -

40. Silver, chrome, stainless steel plating Electric current is used (electrolytic) Nonspontaneous reaction Result: cover a surface with metal plating (spoon, car bumper, etc.)

41. Cathode: reduction (object being plated) Anode: oxidation (metal using to plate) Anode (+) Oxidation Ag o --> Ag + + e - Cathode (-) Reduction Ag + + e - --> Ag o