1. Utilizes relationship between chemical potential energy & electrical energy

2. Redox Reactions battery to start car prevent corrosion
cleaning with bleach (oxidizing agent)
Na, Al, Cl prepared or purified by redox reactions
breathing
O2  H2O and CO2

3. Redox Reactions Synthesis Decomposition Single Replacement Redox rxns
DR rxns NOT redox rxn!
Redox rxns

4. Predicting Redox Reactions
Table J: predict if given redox reaction will occur
metals donate electrons to ion of metals below itself
nonmetal steals electrons from ion of nonmetal below itself

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

6. Predicting Redox Reactions
A + BX  B + AX
If metal A above metal B (Table J): reaction is spontaneous
X + AY  Y + AX
If nonmetal X above nonmetal Y (Table J): reaction is spontaneous

7. Spontaneous or not? Yes Cs + CuCl2  I2 + NaCl  Yes Cl2 + KBr 
Li + AlCl3 
Cs + CuCl2 
I2 + NaCl 
Cl2 + KBr 
Fe + CaBr2 
Mg + Sr(NO3)2 
F2 + MgCl2 
Yes
Yes No
Yes No No
Yes

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

9. Overview of Electrochemistry
TWO kinds of cells:
1. Galvanic or Voltaic (NYS – Electrochemical)
Use spontaneous rxn to produce flow of electrons (electricity) = Exothermic
2. Electrolytic
Use flow of electrons (electricity) to force nonspontaneous rxn to occur = Endothermic

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

11. Electrochemical Cells
spontaneous SR redox rxn: produces flow of electrons
Electrons flow from oxidized substance to reduced substance
Names: Galvanic cells, voltaic cells, or electrochemical cells (NYS)

12. Electrochemical Cells
Redox rxn arranged so electrons forced to flow through wire
When electrons travel through a wire, can make them do work - light a bulb,ring a buzzer
oxidation & reduction reactions must be separated physically

13. Half-Cell Place where each half-reaction takes place ½ cells:
2 needed for complete redox rxn
connected by wire so electrons flow through
connected by salt bridge to maintain electrical neutrality

14. Schematic of Galvanic/Voltaic Cell

15. Parts of a Galvanic/Voltaic Cell
2 half-cells:
One for oxidation rxn &
One for reduction rxn
Each consists of:
container with aqueous solution & electrode (surface where electron transfer takes place)
Wire connects electrodes
Salt bridge connects solutions

16. How much work can you get out of this reaction?
can measure voltage by allowing electrons to travel through voltmeter
galvanic cell is a battery
not easy battery to transport or use in real-life applications

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

18. 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

19. Anode / Cathode How know which electrode is which?
Table J: predict which electrode anode and which electrode is cathode

20. Anode Anode = Oxidation = Electron Donor
anode is metal higher on Table J

21. Cathode Cathode = Reduction = Electron Acceptor
cathode is metal lower on Table J

22. Zn is above Cu, Zn is anode

23. Direction of Electron Flow (through wire):
Anode → Cathode
Direction of Positive Ion Flow
(salt bridge):
Anode → Cathode

24. Positive & Negative Electrode
Negative electrode (anode):
where electrons originate
here it’s Zn electrode
Positive electrode (cathode):
electrode that attracts electrons
here it’s the Cu electrode

25. Aqueous Solution Solution containing ions of same element as electrode
Cu electrode:
Solution: Cu(NO3)3 or CuSO4
Zn electrode:
Solution: Zn(NO3)2 or ZnSO4

26. Salt Bridge migration of ions between half-cells
necessary to maintain electrical neutrality
reaction can not proceed without salt bridge

27. A(s) + BX(aq)  B(s) + AX(aq)
SR rxn occurs during operation of galvanic/voltaic cell
One electrode gains mass (B) and one electrode dissolves (A)
concen of metal ions ↑ in one soln (making AX) & ↓ in other soln (using up BX)

28. Half-Reactions Zn  Zn+2 + 2e- Cu+2 + 2e-  Cu
_________________________
Zn + Cu+2  Zn Cu
Which electrode is dissolving?
Which species is increasing its mass? Zn
Zn+2

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

30. When the reaction reaches equilibrium
voltage is 0!
electrons no longer flow

31. 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: where electrons originate Positive electrode: attracts electrons

32. (-)  Electron flow  wire Positive ion flow  Pb = cathode Al =
anode
Salt bridge
(-)

Pb+2 & NO3-1
Al+3 & NO3-1

33. What are half-reactions?
Al  Al e-
Pb e-  Pb
Al metal is electrode that’s dissolving
Al+3 ions go into solution
Pb+2 ions are in the solution
Ions pick up 2 electrons & plate together on surface of Pb electrode as Pb0

34. Overall Rxn 2(Al  Al+3 + 3e-) + 3(Pb+2 + 2e-  Pb)
_____________________________
2Al + 3Pb+2 + 6e- 2Al+3 + 3Pb + 6e-
2Al + 3Pb+2  2Al+3 + 3Pb

35. 2Al + 3Pb+2  2Al+3 + 3Pb Al Pb Increasing Decreasing
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

36. Application: Batteries

37. Dry Cell

38. Mercury battery

39. Application: Corrosion

40. Corrosion Prevention

41. What’s wrong with this picture?