1. Chemistry Oxidation-Reduction(Redox)

2. I. Introduction Acids donate _____ and bases accept ____ H+H+H+H+H+H+H+H+ proton(s)proton(s) proton(s)proton(s) In other reactions, substances donate or accept electrons The flow of electrons from one substance to the next is how batteries “make” electricity

3. II. Oxidation Numbers Assigning oxidation numbers to substances enable chemists to keep track of the “transfer” of electrons Review: Ionic compounds: give and take e’s Covalent compounds: share e’s

4. In ionic compounds – e’s are really transferred In covalent compounds – e’s are shared unevenly but we pretend that the most electronegative atom gets all the electrons Oxidation Number: The charge that an atom in a molecule would develop if the most electronegative atom in the molecule took the shared electrons (if the e transfer was complete)

5. Let’s calculate ox #’s for HCl… Chlorine is more electronegative We pretend it takes all the e’s That leaves H’s ox # as +1 Cl ox # as -1 Cl ox # as -1

6. III. Rules for Assigning Oxidation Numbers 1. The oxidation number of any pure element is zero. Thus the oxidation number of H in H 2 is zero. 2. The oxidation number of a monatomic ion is equal to its charge. Thus the oxidation number of Cl in the Cl - ion is -1, that for Mg in the Mg 2+ ion is +2, and that for oxygen in O 2- ion is -2. 3. The sum of the oxidation numbers in a compound is zero if neutral, or equal to the charge if an ion.

7. 4. The oxidation number of alkali metals in compounds is +1, and that of alkaline earths in compounds is +2. The oxidation number of F is -1 in all its compounds. 5. The oxidation number of H is +1 in most compounds. Exceptions are H 2 (where H = 0) and the ionic hydrides, such as NaH (where H = -1). 6. The oxidation number of oxygen (O) is -2 in most compounds. Exceptions are O 2 (where O = 0) and peroxides, such as H 2 O 2 or Na 2 O 2, where O = -1. For other elements, you can usually use If no other rules apply, assume ON is the same as the charge taken on in an ionic compound (“the charge it would like to be) For other elements, you can usually use If no other rules apply, assume ON is the same as the charge taken on in an ionic compound (“the charge it would like to be)

8. Examples 1. What are the following ON’s? A. Iron in Fe? 0 B. Iron in Fe 3+? +3 C. Br in MgBr 2 ? D. Br in HBrO 3 +5

9. 2. What is the ON of each atom in the following? A. H 2 0 B. Ca(OH) 2 Ca: +2O: -2 H: +1 C. F 2 CO F: -1C: +4 O: -2 D. NaSO 4 - Na: +1S: +6 O: -2

10. E. C 2 H 4 C: -2 H: +1 F. NCl 3 N: +3 Cl: -1 G. MgCoBr 4 Mg: +2 Co: +2 Br: -1

11. OYO’s 16.1 What is the oxidation number of N in each of the following substances? A. N B. N 3- C. NO 3 - 16.2 What are the oxidation numbers of all atoms in the following compounds? A. LiNH 2 B. N 2 H 2 C. Ca(NO 2 ) 2 D. CO 2 E. BF 4 - F. PO 4 3- G. ClNO H. S 8

12. III. Oxidation and Reduction Oxidation: The process by which an atom loses an electron or electrons Reduction: The process by which an atom gains an electron or electrons LEO goes GER Lose Electrons–Oxidation Gain Electrons- Reduction OIL RIG Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)

13. Examples 1. S atom goes from -2 to +6. Oxidized or reduced? How many electrons? 2. C atom goes from -2 to -4. Oxidized or reduced? How many electrons? 3. An atom goes from +5 to +3. Oxidized or reduced? How many electrons? 4. An atom goes from -6 to -1. Oxidized or reduced? How many electrons?

14. IV. Recognizing Redox Reactions Mg(s) + 2HCl(aq) → H 2 (g) + MgCl 2 (aq) 0 +1 -1 0 +1 -1 0 +1 -1 0 +1 -1 Mg is oxidized H is reduced NaOH(aq) + HCl(aq) → H 2 O(l) +NaCl(aq) +1 -2 +1 +1 -1+1 -2 +1 -1 NOT a redox reaction!

15. 2Mg(s) + O 2 (g) → 2MgO(s) 0 0 +2 -2 0 0 +2 -2 Mg is oxidized O is reduced

16. Examples – Redox, or not? A. Ca(s) + Cl 2 (g) → CaCl 2 (s) 0 0 +2 -1 0 0 +2 -1 Ca oxidized Cl reduced B. Ag + (aq) + OH - (aq) → AgOH(s) +1 -2 +1 +1 -2 +1 +1 -2 +1 +1 -2 +1 not a redox C. 3IF 5 (aq) + 2Fe(s) → 2FeF 3 (aq) + 3IF 3 (aq) +5 -1 0 +3 -1+3 -1 +5 -1 0 +3 -1+3 -1 Fe oxidized I reduced

17. Thionin – Two-Faced Solution Thio + + 2Fe 2+ + 2H + ↔ ThioH 2 + + 2Fe 3+ Purple colorless light activated light activated Methylene Blue O 2 (g) → O 2 (dissolved) O 2 (dissolved) + methylene blue → methylene blue colorless blue colorless blue Glucose + OH - → glucoside Glucoside + methylene blue → methylene blue + OH - blue colorless

18. ** When a redox reaction occurs, there will always be atoms that are oxidized and atoms that are reduced The thing that is oxidized is the reducing agent (provides electrons) The thing that is reduced is the oxidizing agent (takes electrons)

19. OYO’s 16.5 For each of the following, determine whether or not it is a redox reaction. If it is a redox, identify what was oxidized, what was reduced, & the oxidizing and reducing agents. A. Cu 2+ (aq) + Zn(s) → Zn 2+ (aq) + Cu(s) B. 2H 2 O → 2H 2 + O 2 C. NaCl + AgF → AgCl + NaF D. Ca(s)+ 2HNO 3 (aq) → Ca(NO 3 ) 2 (aq)+ H 2 (g)

20. IV. How Batteries Work The fact that atoms can lose and gain electrons in a reaction is the bases for how a battery works Some reactions (like acid-base) involve the rearrangement of electrons Others (like formation, decomposition, combustion and displacement reactions) involve the transfer of electrons (redox) Displacement reactions: an ion (or atom) in a compound is replaced by an ion (or atom) of another element

21. CuSO 4 (aq) + Zn(s) → ZnSO 4 (aq) + Cu(s) Cu 2+ (aq) + Zn(s) → Zn 2+ (aq) + Cu(s) +2 0+2 0 +2 0+2 0

22. How could you get the electrons that Zn loses to Cu to create a current? You’d have to connect them!

23. A simple battery “salt bridge” – Permits the migration of ions.These keep the Zn 2+ solution from becoming too positive (as negative electrons leave), and the other solution from becoming too negative (as negative electrons arrive) -- porous barrier: does not allow solutions to mix, only permits ions to flow to both solutions

24. Anode and Cathode Anode: where oxidation occurs -- negative in galvanic cells Cathode: where reduction occurs -- positive in galvanic cells

25. As electrons travel through the wire, electricity is created (This schematic is reversed from the one before.)

27. Examples Draw a diagram of the Galvanic cell, indicating the electron flow, labeling the anode and cathode, and indicating the positive and negative sides of the Galvanic cell for the following reactions: A. 2Ag + (aq) + Mg(s) → 2Ag(s) + Mg 2+ (aq) B. I 2 (aq) + Fe(s) → 2I - (aq) + Fe 2+

28. OYO’s Draw a diagram of the Galvanic cell, indicating the electron flow, labeling the anode and cathode, and indicating the positive and negative sides of the Galvanic cell for the following reactions: 16.6 2H + (aq) + Mn(s) → H 2 (g) + Mn 2+ (aq) 16.7 3Cl 2 (aq) + 2Al(s) → 2Al 3+ (aq) + 6Cl - (aq)

29. V. Real Batteries A. The Dry Cell… (no aqueous substances) Anode:Zn(s) → Zn 2+ (aq) + 2e - Cathode: 0 +2 2NH 4 + (aq) + 2MnO 2 (s) → Mn 2 O 3 (s) + 2NH 3 (aq) + H 2 O(l) +4 +3 +4 +3

30. B. The Mercury Battery Anode:Zn(Hg) + 2OH - (aq) → ZnO(s) + H 2 O(l) + 2e- 0 +2 0 +2 Cathode: HgO(s) + H 2 O(l) + 2e - → Hg(l) + 2OH - (aq) +2 0 +2 0

31. Why do batteries stop working? One of the reactants is used up. The salt bridge degrades.

32. Rechargeable batteries They “charge” by reversing the reactions to re-create the original substances Eventually stop working because the salt bridge degrades

33. B. The Lead Acid Battery Anode: Pb(s) + SO 4 2- (aq) → PbSO 4 (s) + 2e - 0 +2 0 +2Cathode: PbO 2 (s) + 4H + (aq) + SO 4 2- (aq) + 2e - → PbSO 4 (s) + 2H 2 O(l) +4 +2

35. Lead-acid batteries are charged by your car’s alternator. Stop working because some of the PbSO 4 (s) falls to the bottom of the battery and is removed from the reaction

36. VI. Corrosion All redox reactions are not constructive!! 2Fe(s) + 2H 2 O(l) + O 2 (g) → 2Fe(OH) 2 (s) 0 0 +2 -2 0 0 +2 -2 (Fe(OH) 2 is further oxidized to Fe 2 O 3 (rust) Corrosion costs US Govt $276 Billion/year

37. OYO 16.8 Consider the lead acid battery: Pb(s) + PbO 2 (s) + 4H + (aq) + 2SO 4 2- (aq) → 2PbSO 4 (s) + 2H 2 O(l) What is being oxidized and what is being reduced?