1. Chp 17 Electrochemistry

2. Galvanic Cells
A device in which chemical energy is converted into electrical energy
Oxidation – reduction (redox) reactions
e- ‘s transfer from reducing agent to oxidizing agent

3. Galvanic Cell e- e- e- Salt Bridge Why? Zn Cu Zn 2+ Cu 2+ SO4 2-
Voltage meter e- e- e-
Salt
Bridge
Why? Zn Cu
Zn 2+
SO4 2-
Cu 2+
SO4 2-

4. What is happening to the Zn and the Cu?
Zn is making electrons (oxidization) Cu is gaining electrons (reduction) Anode where electrons are produced. Cathode where electrons are used. Cell potential cell (volt) or (V)
Zn disappearing
Forming Cu

5. Standard Reduction Potentials
Anode: Zn(s)  Zn 2+(aq) + 2 e- Cathode: Cu 2+(aq) + 2 e-  Cu (s) ocell = ocat + oan ocell = (+0.76V) + (+0.34V) = V Must be positive

6. Examples
Fe 3+(aq) + Cu (s)  Cu 2+(aq) + Fe 2+(aq) Half reactions Red: ( Fe e -  Fe 2+) Ox: Cu  Cu e – 2Fe 3+(aq) + Cu (s)  Cu 2+(aq) + 2Fe 2+(aq) ocell = ocat + oan ocell = 0.43 V
o = 0.77 V 2
o = V
Opposite sign because of Ox
= 0.77 V + (-0.34 V)

7. Line Notation
Zn (s) + Cu 2+(aq)  Zn 2+(aq) + Cu (s) Zn(s) l Zn 2+(aq) ll Cu 2+(aq) l Cu(s)
electrode
electrode