Chapter 19: Electrochemistry

1. Chemistry: The Molecular Science Moore, Stanitski and Jurs Chapter 19: Electrochemistry

2. Electrochemistry • Electrochemistry is the study of the relationship between electricity (e- flow) and chemical changes (redox reactions). • electrochemical cells – electrons from a product-favored redox reaction are harnessed in an electrical current • corrosion of metals - product-favored redox reaction • electrolysis – reactant-favored reaction used in production and purification of many metals and in electroplating processes

3. Oxidation-Reduction Reactions • Oxidation-reduction reactions (redox)involve the transfer of electrons from one species to another. • Oxidationis defined as the loss of electrons. (LEO) • Increase in oxidation number. • Reductionis defined as the gain of electrons. (GER) • Decrease (reduction) in oxidation number. • Oxidation and reduction always occur simultaneously.

4. Oxidation-Reduction Reactions Consider the following reaction: Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) • The net ionic equation shows the reaction of zinc metal with Cu2+ to produce Zn2+ ion and copper metal. Loss of 2 e- oxidation Gain of 2 e- reduction

5. Oxidation-Reduction Reactions • Redox – recognize by oxidizing/reducing agent, uncombined element as reactant, change in oxidation number. • An oxidizing agent is a species that oxidizes another species; (O2, F2, HNO3) • A reducing agent is a species that reduces another species; (H2, Na) Loss of 2 e-oxidation reducing agent oxidizing agent Gain of 2 e-reduction

6. Oxidation-Reduction Reactions • The concept of oxidation numbers is a simple way of keeping track of electrons in a reaction. • The oxidation number (or oxidation state) of an atom in a substance is the actual charge of the atom if it exists as a monatomic ion. • Alternatively, it is hypothetical charge assigned to the atom in the substance by simple rules.

7. Oxidation Numbers Oxidation Numbers: 1. For uncombined element, 0 2. Charge on monatomic ion 3. F is always -1 4. Other halogens are -1 except when combined with another halogen 5. H is +1 except in metal hydrides (H is -1) 6. O is -2 except when combined with F, or in a peroxide (O2-2) 7. Sum of oxidation states equals charge on ion or molecule

8. Monatomic Ion Charges e- loss = Group # e- gained, Group #-8 +1 +2 +3 -4 -3 -2 -1

9. Oxidation numbers What is the oxidation number of underlined element in the following species: KCrO2 PH4+ CrO42- P4O10 S4O62- Cl2 (+1) + X + 2(-2) = 0 X= +3 X+ 4(+1) = +1 X = -3 X + 4(-2) = -2 X = +6 4X + 10(-2) = 0 4X = +20 X = +5 4x + 6(-2) = -2 4x = +10 2X = 0 X = 0

10. Half-Reactions oxidation half-reaction reduction half-reaction net reaction No electrons appear in the equation for the net reaction. The number of electrons produced by the oxidation reaction is equal to the number of electrons gained in the reduction reaction.

11. Redox Reactions

12. Example Write oxidation and reduction half reactions for the following net redox equation:

13. Example Write oxidation and reduction half reactions for the following net redox equations:

14. Balancing Redox Equations In acidic solution 1. Assign oxidation numbers +3 +7 +2 +4 2. Write and balance half-reactions To balance hydrogen we can add H+ - acidic solution

15. +3 +7 +2 +4 2. Write and balance half-reactions To balance oxygen we can add water To balance hydrogen we can add H+ - acidic solution

16. 3. Balance charges (electrons) 5 × ( ) 2 × ( ) 4. Add the half-reactions

17. 5. Cancel the equal amounts of reactants and products that appear on both sides 6. Add enough water (on both sides) to convert H+ to H3O+

18. 6. Check atoms and charges Charges: + 4 on both sides C = 10 Atoms: Mn = 2 O = 34 H = 28

19. Balancing Redox Equations In acidic solution 1. Assign oxidation numbers 0 +6 +3 +2 2. Write and balance half-reactions

20. 0 +6 +3 +2 2. Write and balance half-reactions To balance oxygen we can add water To balance hydrogen we can add H+ - acidic solution

21. 3. Balance charges (electrons) 3 × ( ) 4. Add the half-reactions

22. 5. Cancel the equal amounts of reactants and products that appear on both sides 6. Add enough water (on both sides) to convert H+ to H3O+

23. 6. Check atoms and charges Charges: + 12 on both sides Cr = 2 Atoms: Zn = 3 O = 21 H = 42

24. Balancing Redox Equations In basic solution 0 +3 +2 +2 1. Assign oxidation numbers 2. Write and balance half-reactions To balance oxygen we can add water To balance hydrogen we can add H+ (we will add OH- later)

25. 0 +3 +2 +2 2. Write and balance half-reactions To balance oxygen we can add water To balance hydrogen we can add H+ (we will add OH- later)

26. 3. Balance charges (electrons) charges are balanced 4. Add enough OH- to react with H+ to produce H2O 5. Add the two half-cell reactions

27. 3 6. Cancel equal amounts of reactants or products that appear on both sides

28. 6. Check atoms and charges Charges: 0 on both sides Cd = 1 Atoms: Ni = 2 O = 6 H = 6

29. Balancing Redox Equations In basic solution In basic solution, aluminum metal forms Al(OH)4- ion as it reduces NO3- ion to ammonia. Write a balanced equation for this reaction

30. Balancing Redox Equations In basic solution -3 0 +5 +3 1. Assign oxidation numbers 2. Write and balance half-reactions To balance oxygen we can add water To balance hydrogen we can add H+ (we will add OH- later)

31. -3 0 +5 +3 2. Write and balance half-reactions To balance oxygen we can add water To balance hydrogen we can add H+ (we will add OH- later)

32. 3. Balance charges (electrons) 8 ×( ) 3 × ( )

33. 4. Add enough OH- (to both sides) to react with H+ to produce H2O 5. Cancel equal amounts of products and reactants that appear on both sides 6. Add the reactions

34. 6. Check atoms and charges Charges: -8 on both sides Al = 8 Atoms: N = 3 O = 32 H = 41

35. Practice • Balance the following reactions in acidic solution: • HNO3 + H3PO3 NO + H3PO4 • 2. VO+2 + Sn2+ V+3 + Sn4+ • Balance the following reactions in basic solution: • PbO2 + Cl- ClO- + Pb(OH)2 • 2. Cu+2 + H2 Cu + H2O

36. Electrochemical Cells • Anode • Cathode • Electron flow • Salt bridge Negative counter ion would build up without the salt bridge.

37. Electrochemical cells • Also called voltaic cells (Allesandro Volta constructed the first electrochemical cell in 1800) or batteries • redox reactions are taking place in half-cells • zinc and copper strips are called electrodes • anode (-) – the electrode where oxidation (loss of electrons) occurs. (two vowels) • cathode (+) - the electrode where reduction (gain of electrons) occurs. (two constantans) • The electrons that are lost by zinc (anode) pass through the wire and a light bulb to the copper metal (cathode).

38. Electrochemical cells • salt bridge –porous barrier with solution of salt that allows ions to flow between the electrode compartments. The flow of ions completes the electrical circuit, allowing current to flow. If the salt bridge is removed from this battery, the flow of electrons will stop. Shorthand notation for electrochemical cells: Zn(s)| Zn2+(aq) || Cu2+(aq)|Cu(s) anode electrode cathode electrode salt bridge

39. Electrochemical cells • Draw a diagram for a salt bridge cell for the following reaction. • Label anode and cathode and indicate the direction of current flow through-out the circuit. • Write the two half-cell reactions. • Write a shorthand notation for the cell. • Cd (s) + Co+2(aq) Cd+2(aq) + Co (s)

40. Cell Voltage Standard voltage (Eº) – measured at standard conditions all reactants and products must be present as pure solids, liquids, gases (at 1 bar pressure) or solutes at 1M concentration temperature is 25ºC (298K) E°cathode and E°anode are values for the half-reactions written as reductions. E°red and E°ox are the values for half-reactions written as a reduction and oxidation, respectively. If E°cell is positive, a current is produced (product favored or spontaneous).

41. Cell Voltage Since we can only measure the differences in potential energy, the hydrogen electrode is chosen as a standardreference (Eº = 0.0V) and all other electrodes are compared to it. SHE – Standard Hydrogen Electrode (1 bar and 1M) 2H3O+(aq) + 2e- H2(g) + 2H2O(l) Eº = 0.0V

42. easy to reduce (good oxidizing agent) standard hydrogen electrode hard to reduce, easy to oxidize (good reducing agent)

43. Zn(s)  Zn2+ +2e- E º ox = ? 2H3O+(aq) + 2e- H2(g) + 2H2O(l) E º red = 0 Notice that in table 19.1 we have -0.76, since all values listed are reduction potentials.

44. Cu2+ +2e- Cu(s) E ºred = ? H2(g) + 2H2O(l)  2H3O+(aq) + 2e- E º ox = 0

46. Electrochemical cells • Draw a diagram for a salt bridge cell for the following reaction. • Label anode and cathode and indicate the direction of current flow through-out the circuit. • Write the two half-cell reactions. • Write a shorthand notation for the cell. • Mg (s) + 2H+1(aq) Mg+2(aq) + H2(g)

47. ExampleStandard Reduction Potentials Will aluminum metal react with tin(IV) solution? If yes, what is Eº for the reaction. Sn4+(aq) + 2e- Sn2+(aq) Eº = + 0.15 Al3+(aq) + 3e-  Al(s) Eº = - 1.66 Sn4+(aq) is above Al+3 in the table. Sn+4 is more easily reduced. Therefore, it will oxidize Al(s)

48. Sn4+(aq) + 2e- Sn2+(aq) Eº = + 0.15 Al3+(aq) + 3e-  Al(s) Eº = - 1.66 cathode (reduction) Sn4+(aq) + 2e- Sn2+(aq) Eº = + 0.15 anode (oxidation) Al(s)  Al3+(aq) + 3e- Eº = +1.66 2Al(s) + 3Sn4+(aq)  2Al3+(aq) + 3Sn2+(aq) Do not multiply the E° by 3. 3 ×[ ] 2 × [ ] Notice that the sign is changed since the reaction is reversed.

49. ExampleStandard Reduction Potentials Will copper metal react with Cd2+ solution? If yes, what is Eº for the reaction. Cu2+(aq) + 2e- Cu(s) Eº = + 0.34 Cd2+(aq) + 2e-  Cd(s) Eº = - 0.40 Cd2+(aq) is below Cu(s) in the table. It will not oxidize Cu(s) Lets check this with reactions.

50. Cu2+(aq) + 2e- Cu(s) Eº = + 0.34 Cd2+(aq) + 2e-  Cd(s) Eº = - 0.40 cathode (reduction) Cd2+(aq) + 2e- Cd(s) Eº = - 0.40 anode (oxidation) Cu(s)  Cu2+(aq) + 2e- Eº = - 0.34 Cu(s) + Cd2+(aq)  Cd(s) + Cu2+(aq) Notice that the sign is changed since the reaction is reversed The “-” sign indicates that the reaction is reactant favored, the reverse reaction will be more favorable (cadmium reducing copper)