Utilizes relationship between chemical potential energy & electrical energy

1. Electrochemistry Utilizes relationship between chemical potential energy & electrical energy

2. Redox Reactions • Need battery to start car • Prevent corrosion • Bleach is an oxidizing agent • Na, Al, Cl prepared or purified by redox reactions • Breathing • O2 H2O and CO2

3. Redox Reactions • Synthesis • Decomposition • Single Replacement • Double Replacement only is not redox Redox

4. Predicting Redox Reactions • Use Table J to predict if a given redox reaction will occur. • Any metal will donate its electrons to the ion of any metal below it. • Any nonmetal will steal electrons from the ion of any nonmetal below it. Memory Jogger

5. Predicting Single Replacement Redox Reactions • Element + Compound  New Element + New Compound • If the element is above the swapable ion, the reaction is spontaneous. • If the element is below the swapable ion, the reaction is not spontaneous. Memory Jogger

6. Predicting Redox Reactions A + BX  B + AX A & B are metals. If metal A is above metal B in Table J, the reaction is spontaneous. X + AY  Y + AX • X & Y are nonmetals. If nonmetal X is above nonmetal Y in Table J, the reaction is spontaneous. Memory Jogger

7. Which are spontaneous? Yes • Li + AlCl3 • Cs + CuCl2  • I2 + NaCl  • Cl2 + KBr  • Fe + CaBr2  • Mg + Sr(NO3)2  • F2 + MgCl2  Yes No Yes No No Yes

8. Started with Zn(NO3)2 & Cu and AgNO3 & Cu. Which beaker had the Zn ions & which had the Ag ions?

9. Overview of Electrochemistry • TWO kinds of cells (kind of opposites): • Galvanic or Voltaic (NYS – Electrochemical) • Use a spontaneous reaction to produce a flow of electrons (electricity). Exothermic. • Electrolytic • Use a flow of electrons (electricity) to force a nonspontaneous reaction to occur. Endothermic.

10. Redox Half-reaction Oxidation Reduction Cell Half-Cell Electrode Anode Cathode Galvanic Voltaic Electrochemical Electrolytic Salt bridge Vocabulary

11. Electrochemical Cells • Use a spontaneous single replacement redox reaction to produce a flow of electrons. • Electrons flow from oxidized substance to reduced substance. • Called: Galvanic cells, voltaic cells, or electrochemical cells (NYS)

12. Electrochemical Cells • Redox reaction is arranged so the electrons are forced to flow through a wire. • When the electrons travel through a wire, we can make them do work, like light a bulb or ring a buzzer. • So the oxidation & reduction reactions have to be separated physically. OJ clock

13. Al / CuCl2 Lab • Was a redox reaction. • Did NOT force electrons to travel through a wire. Got NO useful work out of system. • Have to be clever in how we arrange things.

14. 2Al + 3Cu+2 2Al+3 + 3Cu Got no useful work because half-reactions weren’t separated.

15. Half-Cell • Where each of the half-reactions takes place. • Need 2 half-cells to have a complete redox reaction. • Need to be connected by a wire for the electrons to flow through. • Need to be connected by a salt bridge to maintain electrical neutrality.

16. Schematic of Galvanic Cell

17. Parts of a Voltaic Cell • 2 half-cells: oxidation & reduction • Each half-cell consists of a container of an aqueous solution & an electrode or surface at which the electron transfer takes place. • Wire connecting electrodes. • Salt bridge connects solutions.

18. How much work can you get out of this reaction? • You can measure the voltage by making the electrons travel through a voltmeter. • The galvanic cell is a battery. Of course, it’s not a very easy battery to transport or use in real-life applications.

19. Electrode Surface at which oxidation or reduction half-reaction occurs. Anode & Cathode

20. An Ox Ate a Red Cat • Anode – Oxidation • The anode = location for the oxidation half-reaction. • Reduction – Cathode • The cathode = location for the reduction half-reaction.

21. Anode / Cathode • How do you know which electrode is which? • Use Table J to predict which electrode is the anode and which electrode is the cathode.

22. Anode • Anode = Oxidation = Electron Donor • The anode is the metal that’s higher in Table J.

23. Cathode • Cathode = Reduction = Electron Acceptor • The cathode is the metal that’s lower in Table J.

24. Zn is above Cu, Zn is anode http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

25. Notation for Cells ZnZn+2Cu+2Cu

26. Direction of Electron Flow(wire) Anode to Cathode Direction of Positive Ion Flow (salt bridge) Anode to Cathode

27. Positive & Negative Electrode • Negative electrode is where electrons originate – here it’s the Zn electrode. • Positive electrode is electrode that attracts electrons – here it’s the Cu electrode.

28. Aqueous Solution • Solution containing ions of the same element as the electrode. • Cu electrode: solution may be Cu(NO3)3 or CuSO4. • Zn electrode: solution may be Zn(NO3)2 or ZnSO4.

29. Salt Bridge • Allows for migration of ions between half-cells. • Necessary to maintain electrical neutrality. • Reaction will not proceed without salt bridge.

30. A(s) + BX(aq)  B(s) + AX(aq) • Single replacement rxn occurs during operation of galvanic cell. • One electrode will gain mass (B) and one electrode will dissolve (A). • The concentration of metal ions will increase in one solution (making AX) & decrease in one solution (using up BX).

31. Half-Reactions Zn  Zn+2 + 2e- Cu+2 + 2e-  Cu _________________________ Zn + Cu+2 Zn+2 + Cu Which electrode is dissolving? Which species is getting more concentrated? Zn Zn+2

32. Zn + Cu+2 Zn+2 + Cu Cu • Which electrode is gaining mass? • Which species is getting more dilute? Cu+2

33. When the reaction reaches equilibrium • The voltage goes to 0.

34. Construct Galvanic Cell with Al & Pb • Use Table J to identify anode & cathode. • Draw Cell, put in electrodes & solutions • Label anode, cathode, direction of electron flow in wire, direction of positive ion flow in salt bridge, positive electrode, negative electrode. • Negative electrode is where electrons originate. Positive electrode attracts electrons.

35. Electron flow  wire Positive ion flow  Pb = cathode Al = anode Salt bridge -  Pb+2 & NO3-1 Al+3 & NO3-1

36. What are half-reactions? Al  Al+3 + 3e- Pb+2 + 2e-  Pb Al metal is the electrode – it’s dissolving. Al+3 ions go into the solution. Pb+2 ions are in the solution. They pick up 2 electrons at the surface of the Pb electrode & plate out.

37. Overall Rxn 2(Al  Al+3 + 3e-) 3(Pb+2 + 2e-  Pb) _____________________________ 2Al + 3Pb+2 2Al+3 + 3Pb

38. 2Al + 3Pb+2 2Al+3 + 3Pb Al • Which electrode is losing mass? • Which electrode is gaining mass? • What’s happening to the [Al+3]? • What’s happening to the [Pb+2]? Pb Increasing Decreasing

39. Application: Batteries

40. Dry Cell

41. Mercury battery

42. Fuel Cell • Converts chemical to electrical energy • Use oxygen or other oxidizing agent • Constant supply of fuel (hydrogen or other hydrocarbon)

43. Application: Corrosion

44. Corrosion Prevention

45. What’s wrong with this picture?

46. Daniell Cell • Invented in 1836 by John Frederic Daniell • Improved battery technology (voltaic pile –problem with hydrogen bubbling) • Definition of the volt