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
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
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
electrons flow from anode to cathode
anode is oxidized
cathode is reduced
negative ions flow from cathode to anode to
relieve neg charge buildup
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
Reduction half reaction Potential , Eo /V
Li+ + e- --> Li(s)-3.045
K+ + e- --> K(s) -2.924
Ca2+ + 2e- --> Ca(s)
-2.76
Na+ + e- --> Na(s)
-2.7109
Al3+ + 3e- --> Al(s)
-1.706
Mn2+ + 2e- --> Mn(s) -1.029
Cd(OH)2(s) + 2e- --> Cd(s) + 2OH-0.812
Zn2+ + 2e- --> Zn(s)
-0.7628
Cr3+ + 3e- --> Cr(s)
-0.74
Fe2+ + 2e- --> Fe(s)
-0.409
PbSO4(s) + 2e- --> Pb(s) + SO42-0.356
Ni2+ + 2e- --> Ni(s)
-0.23
Sn2+ + 2e- --> Sn(s)
-0.1364
Pb2+ + 2e- --> Pb(s)
-0.1263
2H+ + 2e- --> H2(g)
0.0000000
Sn4+ + 2e- --> Sn2+
+0.15
IO3- + 2H2O +4e- --> IO- + 4OH+0.15
SO42- + 4H+ + 2e- --> H2SO3 + H2O +0.20
Cu2+ + 2e- --> Cu(s)
+0.3419
O2(g) + 2H2O(l) + 4e- --> 4OH- +0.401
IO- + H2O + 2e- --> I- + 2OH- +0.485
NiO2(s) + 2H2O + 2e- --> Ni(OH)2(s) + 2OH+0.49
I2(s) + 2e- --> 2I+0.535
Fe3+ + e- --> Fe2+
+0.770
Ag+ + e- --> Ag(s)
+0.7996
ClO- + H2O(l) + 2e- --> Cl- + 2OH+0.90
NO3- + 4H+ + 3e- --> NO(g) + 2H2O(l) +0.96
Br2(l) + 2e- --> 2Br+1.065
O2(g) + 4H+ + 4e- --> 2H2O(l) +1.229
Cl2 + 2e- --> 2Cl+1.3583
MnO4- + 8H+ + 5e- --> Mn2+ + 4H2O +1.507
Au+ + e- --> Au(s)
+1.68
PbO2(s) + 4H+ + SO42- + 2e- --> PbSO4(s) + 2H2O(l)
F2(g) + 2e- --> 2F+2.87
+1.685
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
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?
this would be iron bar in ironsomethinglike Fe(NO3)2 or FeSO4
this would be a copper bar in
coppersomething like Cu(NO3)2 or
CuSO4
Fe/Fe+2//Cu+2/Cu
now, cell potenial, written as ξ ,
ξred and ξ cell
oxid
ξ
oxid =
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
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)
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
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
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
now, electrolytic cells aren't spontaneous,and
don't create electricity. They use electricity ( go
to powerpoint c21 4th and 7th hour)