1. Chapter 17 Electrochemistry
all based on oxidation and reduction, electrochem.
puts electron movement to work!
electrochemistry is the study of the interchange
of chemical and electrical energy
redox reactions can be split into half reactions
and those reactions can be physically seperated
to allow the flow of electrons to be used for energy

2. remember that reducing agents give up electrons and oxidizing agents take electrons
if force the electrons to travel along a wire to get from one half reaction to another, can use thier energy
Galvanic cell- something that converts chemical energy into electrical energy

3. Galvanic cell, labeled
wire and voltometer, to allow e- to flow and measure that flow in volts
salt bridge- allows ions to flow to replace electrons
cathode, the place where reduction occurs
metal and solution of same metal
anode- place where oxidation occurs

4. negative ions flow from cathode to anode to relieve neg charge buildup
electrons flow from anode to cathode
anode is oxidized
cathode is reduced
negative ions flow from cathode to anode to relieve neg charge buildup

5. standard reduction potentials- table that shows the cell potential for the reduction of many metals and ions
all are written as reductions, so when figuring the overall cell potential, one number will have to have the sign reversed

6. Reduction half reaction Potential , Eo /V
Li+ + e- --> Li(s)
K+ + e- --> K(s)
Ca2+ + 2e- --> Ca(s) -2.76
Na+ + e- --> Na(s)
Al3+ + 3e- --> Al(s)
Mn2+ + 2e- --> Mn(s)
Cd(OH)2(s) + 2e- --> Cd(s) + 2OH
Zn2+ + 2e- --> Zn(s)
Cr3+ + 3e- --> Cr(s) -0.74
Fe2+ + 2e- --> Fe(s)
PbSO4(s) + 2e- --> Pb(s) + SO
Ni2+ + 2e- --> Ni(s) -0.23
Sn2+ + 2e- --> Sn(s)
Pb2+ + 2e- --> Pb(s)
2H+ + 2e- --> H2(g)
Sn4+ + 2e- --> Sn
IO3- + 2H2O +4e- --> IO- + 4OH
SO H+ + 2e- --> H2SO3 + H2O +0.20
Cu2+ + 2e- --> Cu(s)
O2(g) + 2H2O(l) + 4e- --> 4OH
IO- + H2O + 2e- --> I- + 2OH
NiO2(s) + 2H2O + 2e- --> Ni(OH)2(s) + 2OH
I2(s) + 2e- --> 2I
Fe3+ + e- --> Fe
Ag+ + e- --> Ag(s)
ClO- + H2O(l) + 2e- --> Cl- + 2OH
NO3- + 4H+ + 3e- --> NO(g) + 2H2O(l) +0.96
Br2(l) + 2e- --> 2Br
O2(g) + 4H+ + 4e- --> 2H2O(l)
Cl2 + 2e- --> 2Cl
MnO4- + 8H+ + 5e- --> Mn2+ + 4H2O
Au+ + e- --> Au(s) +1.68
PbO2(s) + 4H+ + SO e- --> PbSO4(s) + 2H2O(l)
F2(g) + 2e- --> 2F
Notes:
All ions are 1.0 M aqueous. (g) = gas at 1 atmosphere. (s) = solid and (l) = liquid.

7. so lets take a typical reaction and see if we can figure it out!
Lets look at a reaction between Iron and copper
we know from last chapter that iron is more reactive, so lets write it...
Fe + Cu(NO3)2 Fe(NO3)2 + Cu
when we are writing these out, we often use something called line notation, it shortens up the work

8. What gets oxidized? reduced?
Fe goes to Fe +2, so that is oxidized, and
Cu+2 goes to Cu, so it is reduced but to be sure we will be checking their cell potentials,
Line notation would have you write
Fe/Fe+2//Cu+2/Cu. oxidized first, // to show
salt bridge
now, how would that look on a cell?

9. this would be a copper bar in coppersomething like Cu(NO3)2 or CuSO4
Fe/Fe+2//Cu+2/Cu
this would be a copper bar in coppersomething like Cu(NO3)2 or CuSO4
this would be iron bar in ironsomething-
like Fe(NO3)2 or FeSO4

10. now, cell potenial, written as ξ oxid, ξred and ξ cell
Fe/Fe+2//Cu+2/Cu
now, cell potenial, written as ξ oxid, ξred and ξ cell
ξ oxid =

11. one very common use of the energy
created by these voltaic / galvanic
spontaneous reactions is batteries.
there are 4 basic types
dry cell and alkaline dry cell
lead storage
nickel cadmium
fuel cells

12. dry cell and alkaline dry cell
These are the batteries that run minor electronics, like AA, D, C etc
the reaction is between Zn/Zn+2 and MnO2/NH+4
they run continuously, across a porous paper that acts as the salt bridge. graphite rod acts as the conductor.
These are inexpensive but give little voltage
alkaline dry cell uses MnO2/water, and gives a little more energy, also doesn't decay as fast, so lasts longer but they are more expensive ( still AAA, C, D, think energizer vs generic)

13. Lead storage batteries
these are rechargable, give much more power, and are therefore much more expensive and heavier.
They are several alternating sheets of lead and lead dioxide, with sulfuric acid acting with it as the salt bridge. Lead is both oxidized and reduced, and the energy from your car as it runs recharges it

14. Nickel is reduced and cadmium is oxidized
Nickel Cadmium batteries
also rechargable, these are much smaller, run things like drills, phones etc.
Nickel is reduced and cadmium is oxidized
have some "memory" if not completely discharged, not as cheap as dry cell

15. Fuel cells
mostly in use for rockets, this is a system that carries large quantities of fuel with it. it is the oxidation of H and reduction of O2, fueled constantly from an external source ( on the rocket) and creates water

16. now, electrolytic cells aren't spontaneous,and
don't create electricity. They use electricity ( go to powerpoint c21 4th and 7th hour)