1. Chapter 18 Oxidation–Reduction Reactions and Electrochemistry

2. Chapter 18 Table of Contents 2 18.1Oxidation–Reduction Reactions 18.2 Oxidation States 18.3 Oxidation–Reduction Reactions Between Nonmetals 18.4Balancing Oxidation–Reduction Reactions by the Half-Reaction Method 18.5Electrochemistry: An Introduction 18.6Batteries 18.7Corrosion 18.8Electrolysis

3. Section 18.1 Oxidation–Reduction Equations Return to TOC 3 Oxidation–reduction reaction (redox reaction) – a chemical reaction involving the transfer of electrons.  Oxidation – loss of electrons  Reduction – gain of electrons http://www.youtube.com/watch?v=Ftw7a5ccubs

4. Section 18.1 Oxidation–Reduction Equations Return to TOC Copyright © Cengage Learning. All rights reserved 4 Exercise In the reaction below Sn(II) _____________. Sn 2+ + 2Fe 3+ → Sn 4+ + 2Fe 2+ a)gains electrons b)is reduced c)is oxidized d)is neither oxidized nor reduced

5. Oxidation States Section 18.2 Return to TOC 5 Allow us to keep track of electrons in oxidation– reduction reactions by assigning charges to the various atoms in a compound. Oxidation States for the Transition Metals

6. Oxidation States Section 18.2 Return to TOC Copyright © Cengage Learning. All rights reserved 6 1.Oxidation state of an atom in an elemental state = 0 2.Oxidation state of monatomic ion = charge of the ion 3.Oxygen =  2 in covalent compounds (except in peroxides where it =  1) 4.Hydrogen = +1 in covalent compounds 5.Fluorine =  1 in compounds 6.Sum of oxidation states = 0 in compounds 7.Sum of oxidation states = charge of the ion in ions Rules for Assigning Oxidation States

7. Oxidation States Section 18.2 Return to TOC Copyright © Cengage Learning. All rights reserved 7 Exercise Find the oxidation states for each of the elements in each of the following compounds: K 2 Cr 2 O 7 CO 3 2- MnO 2 PCl 5 SF 4 K = +1; Cr = +6; O = –2 C = +4; O = –2 Mn = +4; O = –2 P = +5; Cl = –1 S = +4; F = –1

8. Oxidation States Section 18.2 Return to TOC Copyright © Cengage Learning. All rights reserved 8 What are the Oxidation Numbers for each element in the following? H 2 O N 2 KMnO 4 CO 2 CH 4 CHCl 3 He Cu Na 2 Cr 2 O 7 +1 for H, -2 for O Zero for N, elemental state +1 for K, -2 for O, +7 for Mn -2 for O, +4 for C +1 for H, -4 for C +1 for H, -1 for Cl, +2 for C Zero for He, elemental state Zero for Cu, elemental state +1 for Na, -2 for O, +6 for Cr 1(+1 K)=+1 4(-2 O)= -8 -7 1(+1 H)=+1 3(-1 Cl)= -3 -2 2(+1 Na)=+2 7(-2 O)= -14 -12

9. Oxidation States Section 18.2 Return to TOC 9 More Practice!

10. Oxidation–Reduction Reactions Between Nonmetals Section 18.3 Return to TOC 10 2Na(s) + Cl 2 (g)  2NaCl(s) Na  oxidized  Na is also called the reducing agent (electron donor). Cl 2  reduced  Cl 2 is also called the oxidizing agent (electron acceptor).

11. Oxidation–Reduction Reactions Between Nonmetals Section 18.3 Return to TOC 11 CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g) C  oxidized  CH 4 is the reducing agent. O 2  reduced  O 2 is the oxidizing agent.

12. Oxidation–Reduction Reactions Between Nonmetals Section 18.3 Return to TOC Copyright © Cengage Learning. All rights reserved 12 Transfer of electrons Transfer may occur to form ions Oxidation – increase in oxidation state (loss of electrons); reducing agent Reduction – decrease in oxidation state (gain of electrons); oxidizing agent Redox Characteristics 0 2+ 1- 2+ 1- 0 Zn(s) + CuCl 2 (aq) ZnCl 2 (aq) + Cu(s) 14 - Reduction Oxidation

13. Oxidation–Reduction Reactions Between Nonmetals Section 18.3 Return to TOC Copyright © Cengage Learning. All rights reserved 13 Concept Check Which of the following are oxidation – reduction reactions? Identify the oxidizing agent and the reducing agent. a)Zn(s) + 2HCl(aq)  ZnCl 2 (aq) + H 2 (g) b)Cr 2 O 7 2- (aq) + 2OH - (aq)  2CrO 4 2- (aq) + H 2 O(l) c)2CuCl(aq)  CuCl 2 (aq) + Cu(s)

14. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 14 Half–Reactions The overall reaction is split into two half–reactions, one involving oxidation and one reduction. Has electrons as reactants or products 8H + + MnO 4 – + 5Fe 2+ → Mn 2+ + 5Fe 3+ + 4H 2 O Reduction: 8H + + MnO 4 – + 5e – → Mn 2+ + 4H 2 O Oxidation: 5Fe 2+ → 5Fe 3+ + 5e –

15. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 15 1.Identify and write the equations for the oxidation and reduction half – reactions. 2.For each half – reaction: A.Balance all the elements except H and O. B.Balance O using H 2 O. C.Balance H using H +. D.Balance the charge using electrons. The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Acidic Solution

16. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 16 3.If necessary, multiply one or both balanced half – reactions by an integer to equalize the number of electrons transferred in the two half – reactions. 4.Add the half – reactions, and cancel identical species. 5.Check that the elements and charges are balanced. The Half–Reaction Method for Balancing Equations for Oxidation– Reduction Reactions Occurring in Acidic Solution

17. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 17 Cr 2 O 7 2- (aq) + SO 3 2- (aq)  Cr 3+ (aq) + SO 4 2- (aq) How can we balance this equation? First Steps:  Separate into half-reactions.  Balance elements except H and O.

18. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 18 Cr 2 O 7 2- (aq)  2Cr 3+ (aq) SO 3 2- (aq)  SO 4 2- (aq) Balance O’s with H 2 O and H’s with H + Method of Half Reactions 14H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 O(aq) H 2 O(l) + SO 3 2- (aq)  SO 4 2- (aq) + 2H + (aq) How many electrons are involved in each half reaction? Balance the charges.

19. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 19 Method of Half Reactions (continued) 6 e- + 14H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 O(aq) H 2 O(l) + SO 3 2- (aq)  SO 4 2- (aq) + 2H + (aq) + 2e- Multiply whole reactions by a whole number to make the number of electrons gained equal the number of electrons lost. 6 e- + 14H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 O(aq) 3(H 2 O(l) + SO 3 2- (aq)  SO 4 2- (aq) + 2H + (aq) + 2e-) Combine half reactions cancelling out those reactants and products that are the same on both sides, especially the electrons.

20. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 20 Final Balanced Equation: Cr 2 O 7 2- + 3SO 3 2- + 8H +  2Cr 3+ + 3SO 4 2- + 4H 2 O Method of Half Reactions (continued) 6e- + 14H + (aq) + Cr 2 O 7 2- (aq)  2Cr 3+ (aq) + 7H 2 O(aq) 3H 2 O(l) + 3SO 3 2- (aq)  3SO 4 2- (aq) + 6H + (aq) + 6e- 4 8

21. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC Copyright © Cengage Learning. All rights reserved 21 Exercise When the reaction Ce 2+ + Co 2+ → Ce 3+ + Co is balanced, the coefficient in front of Ce 2+ is a)0 b)1 c)2 d)3 Ce 2+ → Ce 3+ +1e- 2e- + Co 2+ → Co 2Ce 2+ + Co 2+ → 2Ce 3+ + Co

22. Section 18.4 Balancing Oxidation–Reduction Reactions by the Half-Reaction Method Return to TOC 22 Exercise Balance the following oxidation – reduction reaction that occurs in acidic solution. Br – (aq) + MnO 4 – (aq)  Br 2 (l)+ Mn 2+ (aq) 10Br – (aq) + 16H + (aq) + 2MnO 4 – (aq)  5Br 2 (l)+ 2Mn 2+ (aq) + 8H 2 O(l)

23. Section 18.5 Electrochemistry: An Introduction Return to TOC 23 Electrochemistry The study of the interchange of chemical and electrical energy. Two types of processes:  Production of an electric current from a chemical reaction.  The use of electric current to produce a chemical change.

24. Section 18.5 Electrochemistry: An Introduction Return to TOC 24 Making an Electrochemical Cell 8H + + MnO 4 – + 5e – → Mn 2+ + 4H 2 O Fe 2+ → Fe 3+ + e –

25. Section 18.5 Electrochemistry: An Introduction Return to TOC Copyright © Cengage Learning. All rights reserved 25 Making an Electrochemical Cell If electrons flow through the wire charge builds up. Solutions must be connected to permit ions to flow to balance the charge.

26. Section 18.5 Electrochemistry: An Introduction Return to TOC Copyright © Cengage Learning. All rights reserved 26 Making an Electrochemical Cell A salt bridge or porous disk connects the half cells and allows ions to flow, completing the circuit.

27. Section 18.5 Electrochemistry: An Introduction Return to TOC Copyright © Cengage Learning. All rights reserved 27 Electrochemical Battery (Galvanic Cell) Device powered by an oxidation–reduction reaction where chemical energy is converted to electrical energy. Anode – electrode where oxidation occurs Cathode – electrode where reduction occurs

28. Section 18.5 Electrochemistry: An Introduction Return to TOC Copyright © Cengage Learning. All rights reserved 28 Electrolysis Process where electrical energy is used to produce a chemical change.  Nonspontaneous

29. Section 18.6 Batteries Return to TOC 29 Lead Storage Battery Anode reaction – oxidation Pb + H 2 SO 4  PbSO 4 + 2H + + 2e  Cathode reaction – reduction PbO 2 + H 2 SO 4 + 2e  + 2H +  PbSO 4 + 2H 2 O

30. Section 18.6 Batteries Return to TOC 30 Lead Storage Battery – Overall Reaction Pb (s) + PbO 2 (s) + 2H 2 SO 4 (aq)  2PbSO 4 (s) + 2H 2 O (l) Hydrometer to measure H 2 SO 4 concentration. As the battery discharges the sulfate of the acid precipitates with the lead taking it out of solution and reducing the acid concentration. As the battery is recharged the current goes into dissolving the lead sulfate restoring the acid concentration.

31. Section 18.6 Batteries Return to TOC 31 Electric Potential The “pressure” on electrons to flow from anode to cathode in a battery, like water flow.

32. Section 18.6 Batteries Return to TOC 32 Dry Cell Batteries Do not contain a liquid electrolyte. Acid version Anode reaction – oxidation Zn  Zn 2+ + 2e  Cathode reaction – reduction 2NH 4 + + 2MnO 2 + 2e   Mn 2 O 3 + 2NH 3 + 2H 2 O

33. Section 18.6 Batteries Return to TOC 33 Dry Cell Batteries Do not contain a liquid electrolyte.  Alkaline version – Anode reaction – oxidation Zn + 2OH   ZnO + H 2 O + 2e  – Cathode reaction – reduction 2MnO 2 + H 2 O + 2e   Mn 2 O 3 + 2OH 

34. Section 18.6 Batteries Return to TOC Copyright © Cengage Learning. All rights reserved 34 Dry Cell Batteries Do not contain a liquid electrolyte.  Other Types Silver cell – Zn anode, Ag 2 O cathode Mercury cell – Zn anode, HgO cathode Nickel-cadmium – rechargeable

35. Section 18.7 Corrosion Return to TOC Copyright © Cengage Learning. All rights reserved 35 The oxidation of metals to form mainly oxides and sulfides.  Some metals, such as aluminum, protect themselves with their oxide coating.  Corrosion of iron can be prevented by coatings, by alloying and cathodic protection.  Cathodic protection of an underground pipe.

36. Section 18.8 Electrolysis Return to TOC Copyright © Cengage Learning. All rights reserved 36 Forcing a current through a cell to produce a chemical change that would not otherwise occur.