Electrochemistry
Experiment 12
Oxidation – Reduction Reactions
• Consider the reaction of Copper
wire and AgNO3(aq)
AgNO3(aq)
Cu(s)
Ag(s)
Oxidation – Reduction Reactions
• If you leave the reaction a long time the
solution goes blue!
• The blue is due to Cu2+(aq)
Oxidation-Reduction Reactions
• So when we mix Ag+(aq) with Cu(s) we get Ag(s) and Cu2+(aq)
• Ag+(aq) + 1e-  Ag(s)
• Cu(s)  Cu2+(aq) + 2e-
gain electrons = reduction
lose electrons = oxidation
• The electrons gained by Ag+ must come from the Cu2+
• Can’t have reduction without oxidation (redox)
• Each Cu can reduce 2 Ag+
2Ag+(aq) + 2e-  2Ag(s)
Cu(s)  Cu2+(aq) + 2e2Ag+(aq) + 2e- + Cu(s) 2Ag(s)+ Cu2+(aq) + 2e-
Redox Cu/Ag
E
Cu
electron flow
Ag+
Ag+
Redox Cu/Ag
e = charge on an electron
V = Voltage in a electrochemical cell
E
Cu2
+
ΔE = e.V
Ag
Ag
Redox Reactions & Current
• redox reactions involve the transfer of electrons from
one substance to another
• therefore, redox reactions have the potential to
generate an electric current
• in order to use that current, we need to separate the
place where oxidation is occurring from the place that
reduction is occurring
7
Electric Current Flowing Directly
Between Atoms
8
Electric Current Flowing Indirectly
Between Atoms
Tro, Chemistry: A Molecular Approach
9
Electrochemical Cells
• electrochemistry is the study of redox reactions that
produce or require an electric current
• the conversion between chemical energy and electrical
energy is carried out in an electrochemical cell
• spontaneous redox reactions take place in a voltaic cell
– aka galvanic cells
• nonspontaneous redox reactions can be made to occur in
an electrolytic cell by the addition of electrical energy
Tro, Chemistry: A Molecular Approach
10
Electrochemical Cells
• oxidation and reduction reactions kept separate
– half-cells
• electron flow through a wire along with ion flow through
a solution constitutes an electric circuit
• requires a conductive solid (metal or graphite) electrode
to allow the transfer of electrons
– through external circuit
• ion exchange between the two halves of the system
– electrolyte
Tro, Chemistry: A Molecular Approach
11
Electrodes
• Anode (donates electrons to the cathode)
– electrode where oxidation occurs
– anions attracted to it
– connected to positive end of battery in electrolytic cell
– loses weight in electrolytic cell
• Cathode (attracts electrons from the anode)
– electrode where reduction occurs
– cations attracted to it
– connected to negative end of battery in electrolytic
cell
– gains weight in electrolytic cell
• electrode where plating takes place in
electroplating
Tro, Chemistry: A Molecular Approach
12
Voltaic Cell
the salt bridge is
required to complete
the circuit and
maintain charge
balance
Tro, Chemistry: A Molecular Approach
13
Current and Voltage
• the number of electrons that flow through the system per second
is the current
– unit = Ampere
– 1 A of current = 1 Coulomb of charge flowing by each second
– 1 A = 6.242 x 1018 electrons/second
– Electrode surface area dictates the number of electrons that
can flow
• the difference in potential energy between the reactants and
products is the potential difference (the potential for an electric
field to cause an electrical current)
– unit = Volt
– 1 V of force = 1 J of energy/Coulomb of charge
– the voltage needed to drive electrons through the external
circuit
– amount of force pushing the electrons through the wire is
called the electromotive force, emf
Tro, Chemistry: A Molecular Approach
14
Cell Potential
• the difference in potential energy between the anode the
cathode in a voltaic cell is called the cell potential
• the cell potential depends on the relative ease with which the
oxidizing agent is reduced at the cathode and the reducing
agent is oxidized at the anode
• the cell potential under standard conditions is called the
standard emf, E°cell
– 25°C, 1 atm for gases, 1 M concentration of solution
– sum of the cell potentials for the half-reactions
Tro, Chemistry: A Molecular Approach
15
Standard Reduction Potential
• a half-reaction with a strong tendency to occur has a large +
half-cell potential
• when two half-cells are connected, the electrons will flow
so that the half-reaction with the stronger tendency will
occur
• we cannot measure the absolute tendency of a halfreaction, we can only measure it relative to another halfreaction
• we select as a standard half-reaction the reduction of H+ to
H2 under standard conditions, which we assign a potential
difference = 0 V
– standard hydrogen electrode, SHE
16
Tro, Chemistry: A Molecular Approach
17
Half-Cell Potentials
• SHE reduction potential is defined to be exactly 0 V
• half-reactions with a stronger tendency toward reduction than the SHE have
a + value for E°red
• half-reactions with a stronger tendency toward oxidation than the SHE have
a + value for E°red
• ΔE°cell = E°oxidation + E°reduction
– E°oxidation = -E°reduction
– when adding E° values for the half-cells, do not multiply the half-cell E°
values, even if you need to multiply the half-reactions to balance the
equation
• ΔGocell=-nFΔE°cell
Tro, Chemistry: A Molecular Approach
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Electrochemical Cell Summary
The
potential
bereactants,
calculated
knowing
the
reduction
This
manifests
as acan
ispotential
related
todifference
the free
energy
Ecell
, across
of standard
thethe
reaction
electrodes.
according
Where
to -qEcell
2+
2+, and
The cell
differing
stability
of
(Zn(s),
Cu
(aq)),
and
products
(Zn
potentials.
These
becellused
to when
find
Eoan
for the reaction
at the
cathode,
is
the
change
potential
energy
of
negative
charge
(-q)
the
relation
in
Gacell
=can
-nFE
red amount
Cu(s)),
creates
potential
energy
gradient
through
which
the
charges
o ).
and
Eooxfrom
(= - Ethe
Then
Eoox+ Eored
passes
anode
toEothe
cathode
red
cell
migrate
(from
high
energy
to =low).
2+-(aq)
Zn2+(aq)
--> Zn(s)
Zn(s)
--> +Zn2e
+ 2e-
-0.76VEox= 0.76V
Cu2+(aq) + 2e- --> Cu(s)
Ered=0.34V
Ecell = 0.76V+0.34V = 1.1V
e-
Ecell=1.1 V
salt bridge
cathode
anode
Zn (s)--> Zn2+ (aq)+ 2e-
Cu2+(aq)+ 2e- --> Cu(s)
Tro, Chemistry: A Molecular Approach
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Tonight
• Construction of Voltaic Cells and Measurement of Cell
Potentials
Trial
electrodes
Trial
electrodes
1
Cu/Zn
4
Zn/Pb
2
Cu/Pb
5
Zn/Ni
3
Cu/Ni
6
Pb/Ni
• Use the corresponding 0.1 M metal sulfate of the same metal
as the electrode in the half cell
• Construct a salt bridge
• Measure the voltage