1. Electrochemistry

2. Voltaic Cell (or Galvanic Cell)
The energy released in a spontaneous redox reaction can be used to perform electrical work.
A voltaic cell is a device in which the transfer of electrons takes place through an external pathway rather than directly between reactants.
In a voltaic cell, chemical energy is changed to electrical energy.

3. Half Cells
The voltaic cell is thought of as being comprised of two "half-cells."
One cell is where oxidation occurs and the other is reduction.

4. Half Cells There are two half cells in an electrochemical cell.
Each half cell is made of a metal strip in contact with a solution of its ions.
In the oxidation half cell, the metal electrode loses electrons.
In the reduction half cell, the ions in solution gain electrons and create metal atoms which deposit onto the electrode.

6. Anode/Cathode
The two solid metals that are connected by the external circuit are called electrodes.
The electrode at which oxidation occurs is called the anode. This electrode strongly attracts negative ions in the solution, these ions are then called anions.
The electrode at which reduction occurs is called the cathode. Cations will be deposited onto this electrode as a metal.

7. Half Reactions Anode (oxidation half-reaction):
Zn (s) → Zn+2 (aq) + 2e-
Cathode (reduction half-reaction):
Cu+2 (aq) +2e- → Cu (s)
Determining half reactions will be shown later.

8. OIL – RIG Remember the acronym OIL – RIG
Oxidation Is Loss of electrons
Reduction Is Gain of electrons

9. Red-Ox Reaction
In a red-ox reaction, one substance must be oxidized and another substance must be reduced.
The substance that is oxidized is the reducing agent.
The substance that is reduced is the oxidizing agent.

10. Cell Operation
Electrons become available when the zinc metal is oxidized at the anode.
The electrons flow through the external circuit to the cathode, where they are consumed as Cu+2 is reduced. Because zinc is oxidized in the cell, the zinc electrode loses mass, and the concentration of the Zn+2 solution increases as the cell operates. At the same time, the Cu electrode gains mass, and the Cu+2 solution becomes less concentrated as the Cu+2 is reduced to Cu (s).

11. Cell Operation
As the voltaic cell operates, oxidation of Zn introduces additional Zn+2 ions into the anode compartment. Unless a means is provided to neutralize this positive charge, no further oxidation can take place.
At the same time, the reduction of Cu+2 at the cathode leaves an excess of negative charge in solution in that compartment.

12. Salt Bridge
Electrical neutrality of the system is maintained by a migration of ions through the porous glass disc (salt bridge) that separates the two compartments.
A salt bridge consists of a U-shaped tube that contains an electrolyte solution whose ions will not react with other ions in the cell or with the electrode materials.

13. Salt Bridge
As oxidation and reduction proceed at the electrodes, ions from the salt bridge migrate to neutralize charge in the cell compartments.
The next slide shows the Zn+2 ions migrating into the salt bridge from the left cell and the SO4+2 ions migrating into the salt bridge from the right cell.

14. Salt Bridge
The salt bridge is required for completing the circuit. It allows the movement of ions from one solution to the other so that a charge does not build up in either of the cells.

16. Ion Migration
Anions migrate toward the anode and cations toward the cathode.
No measurable electron flow will occur through the external circuit unless a means is provided for ions to migrate through the solution from one electrode compartment to another, completing the circuit.

17. Electron Flow
In any voltaic cell the electrons flow from the anode through the external circuit to the cathode.
Because the negatively charged electrons flow from the anode to the cathode, the anode in a voltaic cell is labeled with a negative sign and the cathode with a positive sign.

18. Oxidation Numbers Remind yourself how to determine oxidation numbers.
Oxidation numbers show what the charge of each atom would be (in a molecule or ion), if each atom were an ion.
Go to the text and revisit the Rules for Assigning Oxidation Numbers section.

19. Balancing Red-Ox Equations
Obey the Law of Conservation of Mass
Gain and loss of electrons must be balanced

20. Half Reactions Cu + 2Ag+ → 2 Ag + Cu+2
Determine the oxidation number for each substance in the reaction
Write two half reactions based on one substance being oxidized and one being reduced.

21. Half Reactions Cu → Cu+2 + 2e- oxidation 2e- + 2Ag+ → 2 Ag reduction
Half reactions show the number of electrons gained or lost by the substance, standard redox equations do not.
Go to ChemReview packet to practice half reaction writing.

22. Review Sample Ex: 20.4
Try Practice Exercises

23. Cell EMF
What is the driving force that pushes the electrons through an external circuit in a voltaic cell?
An oxidizing agent in one compartment pulls electrons through a wire from a reducing agent in the other compartment.
The pull (or driving force) on the electrons is called the cell potential ( Ecell ) or electromotive force (emf) of the cell.
Or, electrons flow from the anode to the cathode (in a voltaic cell) because of a difference in potential energy.

24. Potential Energy of Electrons
The potential energy of electrons is higher in the anode than in the cathode. Therefore, electrons spontaneously flow through an external circuit from the anode to the cathode.
For any cell reaction that proceeds spontaneously (i.e. voltaic cell), the cell potential will be positive.

25. Standard Conditions Standard Conditions:
* 1M concentrations for reactants and products in solution
* 1 atm pressure (for gases)
* 25 ˚C
Under Standard Conditions the EMF is called the standard emf or the standard cell potential (E ˚cell )
Keep in mind that the superscript ˚ denotes standard-state conditions.

26. Cell Potential
The cell potential of a voltaic cell depends on the particular cathode and anode half-cells involved.
Standard potentials have been assigned to each individual half-cell, and then use the half-cell potentials to determine E ˚cell .
The cell potential is the difference between two electrode potentials (anode and cathode).
The potential associated with each electrode is chosen to be the potential for reduction to occur at that electrode.

27. Standard Reduction Potentials
Standard electrode potentials are tabulated for reduction reactions, so they are called standard reduction potentials ( E ˚red ).
E ˚cell = E ˚red (cathode) - E ˚red (anode) Or
E ˚cell = E ˚red (cathode) + E ˚ox (anode)
The hydrogen half-reaction was used as a reference for all of the other half-reactions.
The standard reduction potential for the hydrogen half-reaction is assigned a standard reduction potential of exactly 0 V.
It is also called the standard hydrogen electrode. (SHE)

28. Calculating EMF
We will use standard reduction potentials to calculate the EMF of a voltaic cell.
Because electrical potential measures potential energy per electrical charge, standard reduction potentials are intensive properties.
Changing the stoichiometric coefficient in a half-reaction does not affect the value of the standard reduction potential.

29. FYI
In general, the more active the metal, the lower its potential.

30. Metal Activity
The most active metal, Li, has the lowest tendency to gain electrons (to be reduced).
Li has the strongest tendency of the metals listed to lose electrons (to be oxidized).

31. Metal Activity
The least active metal, Au, is at the top of the table. It has the greatest tendency to be reduced.
Gold also has the lowest tendency to lose electrons (to be oxidized).

32. A metal that is listed lower in the table of reduction potentials is more active, so it can replace any metal that is higher.
The more active metal will replace one that is less from a solution of its ions.

35. In the cell diagram:
Zinc replaces copper
Zn + CuSO4 → ZnSO4 + Cu

36. Calculating EMF
Try sample exercise 20.5

37. E ˚red and spontaneity
The more positive the E ˚red value for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced, and, therefore, to oxidize another species.
Try sample exercise 20.8

38. EMF and Free Energy Change
ΔG = - nFE
n is a positive # that represents the # of e- transferred in the reaction
F is Faraday’s constant (this constant is the quantity of electrical charge on 1 mole of electrons)
1F = 96,500 C/mol = 96,500 J / V mole
The units for ΔG are J/mole

39. EMF and Free Energy Change
ΔG = - nFE
If n and F are positive values, a positive value of E leads to a negative ΔG.
Remember that a negative ΔG indicates a spontaneous reaction.
The equation can be altered slightly if the reactants and products are in their standard states: ΔG˚ = - nFE˚

40. Concentration and Cell EMF
Remember ΔG = ΔG˚ + RT ln Q
Combining
ΔG = ΔG˚ + RT ln Q and ΔG = - nFE
You have the Nernst equation:
E = E˚ – RT ln Q nF

41. Nernst Equation The Nernst equation can be expressed two ways:
E = E˚ – RT ln Q nF
E = E˚ – 2.303RT log Q

42. Nernst Equation Variation
The equation can be simplified if the cell is run under standard conditions:
E = E˚ – V log Q n
* Temp. is 298 K

43. Use the Nernst Equation to work through the example on the bottom of page 773
Try the sample exercise and practice exercise

44. Electrolysis
This process is the opposite of a voltaic cell.

45. Ch 20 Problems
6, 7, 9, 14, 24, 25 a,b, 27, 29, 31, 36, 43, 44, 46, 49, 51, 52, 56, 59