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 Often Redox Single Replacement
Double Replacement only is not redox
Often Redox
Always 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 Cs + CuCl2  I2 + NaCl  Yes Cl2 + KBr 
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. Vocabulary Redox Half-reaction Oxidation Reduction Cell Half-Cell
Electrode
Anode
Cathode
Galvanic
Voltaic
Electrochemical
Electrolytic
Salt bridge

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 Fat 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

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 Cu
Which electrode is dissolving?
Which species is getting more concentrated? Zn
Zn+2

32. Zn + Cu+2  Zn+2 + Cu Cu Cu+2 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 e-
Pb e-  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 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

39. Application: Batteries

40. Dry Cell

41. Mercury battery

42. Application: Corrosion

43. Corrosion Prevention

44. What’s wrong with this picture?

45. Electric Potential
The “pull” or “driving force” on the electrons caused by the cathode is called electric potential.
Analogous to gravitational potential:
Electrons flow from higher electric potential to lower electric potential
Electric potential, or voltage, units of volts (V)
Current is the movement of electrons: amps (A)

46. Electrode Potentials
The difference in potential between an electrode and its solution is known as electrode potential.
To predict the tendency of a substance to be reduced, a table of reduction potentials has been generated.
Compare to a std hydrogen electrode (SHE) at standard conditions (1.0M, 1atm, 25oC)
Called standard electrode potential, Eo

47. Cell Voltage Eocell = Eocathode – Eoanode
Used to predict if a redox reaction will occur spontaneously  + Eocell
The half-rxn that has the more negative std reduction potential will be the anode (hence, the reason for the negative sign).
To write the half-rxn as oxidation, reverse the sign. This is done automatically in the above equation.

48. Example Problem
Write the overall cell reaction, and calculate the cell potential (Eocell) for a voltaic cell consisting of the following half-cells:
An iron (Fe) electrode in a solution of Fe(NO3)3 and,
A silver (Ag) electrode in a solution of AgNO3.
Note: When a half-rxn is multiplied by a constant, the Eo value is NOT affected.