1. Zn(s) + CuSO4(aq)→ ZnSO4 (aq) + Cu(s)
Write the equation for the reaction
Zn(s) + CuSO4(aq)→ ZnSO4 (aq) + Cu(s)
What type of reaction is this?
single replacement

2. Zn(s) + Cu2+(aq)→ Zn2+(aq) + Cu(s)
Write the net ionic equation for the overall reaction
Zn(s) + Cu2+(aq)→ Zn2+(aq) + Cu(s)
Zinc is being oxidized while the copper is being reduced. Why?
Zn goes from 0 to +2 = oxidized
Cu goes from +2 to 0 = reduced
What is the oxidizing agent? reducing agent?
copper (II) sulfate = oxidizing agent
zinc metal = reducing agent

3. Redox reactions involve electron transfer:
Lose e - = Oxidation
Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s)
Gain e - = Reduction
Cu goes from 0 to +2 = oxidized
Ag goes from +1 to 0 = reduced
silver ion = oxidizing agent
copper metal = reducing agent

4. Standard Cell Potential
Just as the water tends to flow from a higher level to a lower level, electrons also move from a higher “potential” to a lower potential. This potential difference is called the electromotive force (EMF) of cell and is written as E°cell.
The standard for measuring the cell potentials is called a SHE (Standard Hydrogen Electrode).
Description of SHE
Reaction 2H+(aq, 1M)+ 2e- H2(g, 101kPa) E0= 0.00 V

5. oxidizing agents Standard Reduction Potentials
Many different half cells can be paired with the SHE and the standard reduction potentials for each half cell is obtained. Check the table for values of reduction potential for various substances:
Would substances with high reduction potential be strong oxidizing agents or strong reducing agents? Why?
oxidizing agents
The reduction potential tells you the likelihood that a substance will be reduced the stronger the reduction potential the better the substance will serve as an oxidizing agent.

7. Review - Activity Series
For metals, the higher up the chart the element is, the more likely it is to be oxidized. This is because metals like to lose electrons, and the more active a metallic element is, the more easily it can lose them.
For nonmetals, the higher up the chart the element is, the more likely it is to be reduced. This is because nonmetals like to gain electrons, and the more active a nonmetallic element is, the more easily it can gain them.

8. Metal Activity
2K + FeCl2
REACTION
Metallic elements start out with a charge of ZERO, so they can only be oxidized to form (+) ions.
The higher of two metals MUST undergo oxidation in the reaction, or no reaction will happen.
The reaction 2K + FeCl2  2KCl + Fe WILL happen, because K is being oxidized, and that is what the table says should happen.
The reaction Fe + 2KCl  FeCl2 + 2K will NOT happen.
Fe + 2KCl
NO REACTION

9. Voltaic Cells (Galvanic Cells)
A voltaic cell converts chemical energy from a spontaneous redox reaction into electrical energy. Ex: Cu and Zn voltaic cell (More positive reduction potential is the cathode)
Key Words:
Cathode
Anode
Salt Bridge
How a Voltaic Cell Works: An Ox, Red Cat
A reaction is spontaneous if the metal with higher reduction potential is made cathode. . t
Positive terminal of a voltaic cell
Negative terminal of a voltaic cell
Connects the half-cells of voltaic cell to prevent the cell from rapidly running its reaction to equilibrium

10. Voltaic Cells
Produce electrical current using a spontaneous redox reaction
Used to make batteries!
Use Reference Table to determine the metals to use
(-) anode (lower reduction potential)
(+) cathode (higher reduction potential)

11. Making Voltaic Cells Oxidation: Reduction: Zn Zn2+ + 2e-
E0= V
Reduction:
Cu2+ + 2e-  Cu
E0= V

12. Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s)
Cell Diagram Notation
Cell diagram notation is shorthand that expresses a certain reaction in an electrochemical cell.
The half-cells are separated by a double slash, which represent a salt bridge. (anode always on the left)
Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s)
Cell Half Reactions
Cu2+ (aq) + 2e-(aq)→ Cu(s) E°= V
Zn(s) → Zn2+(aq) + 2e-(aq) E°= V
Eocell = V