Electrochemical Cells
Electrochemical and Electrolytic Cells
An electrochemical cell uses chemical reactions to produce
electrical current, to release potential electrical energy.
An electrolytic cell uses electrical energy to produce desired
chemical reactions at electrodes like electroplating an
electrode.
Electrochemical Cell
Electrolytic Cell
Electrochemical and Electrolytic Cells
The anode in an electrochemical cell (galvanic cell) is a
negative electrode or terminal while the cathode is a
positive terminal.
The anode in an electrolytic cell is a positive terminal while
the cathode is a negative terminal.
In both of these cells, the terminal where oxidation occurs is
the anode while the terminal where reduction occurs is the
cathode. This can be remembered with the mneumonic,
“An Ox CaRed”.
Electrochemical Cell
Electrolytic Cell
Parts of an Electrochemical Cell
• The anode is the electrode where oxidation occurs, where
electrons flow away from, where solution anions move
towards.
• The cathode is the electrode where reduction occurs,
where electrons move towards, where solution cations
move towards.
oxidation
reduction
Remembering Electrode Reactions
The memory aid, “an ox cared” , stands for anode-oxidation,
cathode-reduction. This relationship holds true for both
electrochemical cells and electrolytic cells.
Features of an Electrochemical Cell
1.
Referring to a standard
reduction potentials table,
it can be seen that in
comparing the copper and
zinc electrodes, copper
has a greater tendency to
reduce while zinc has a
greater tendency to
oxidize.
Features of an Electrochemical Cell
2. As Zn oxidizes, the Zn+2 ions migrate into solution leaving
electrons behind in the Zn electrode (making it negative).
The Zn electrode is the anode (An ox cared). The anode
loses mass.
3. Copper ions Cu+2 , from the solution next to the copper
electrode, reduce at the copper electrode which plate new
copper metal onto the copper electrode. The Cu electrode
is the cathode. The cathode gains mass.
Features of an Electrochemical Cell
4. Electrons flow from the anode (Zn) to the cathode (Cu),
from – to +.
5. Without the salt bridge, Zn+2 ions would build up around
the Zn electrode pulling onelectrons, keeping them from
flowing away. In addition (without the salt bridge), Cu+2
ions from solution plating on the Cu electrode leave the
solution overall negative (SO4-2 anions from the CuSO4
compound). These anions repel incoming electrons,
stopping the flow.
Features of an Electrochemical Cell
6. The function of the salt bridge is to neutralize the charge
build-up around each electrode due to solution ions.
Positive ions from the salt bridge move to the cathode to
neutralize the minus charge of the solution anion build-up.
Negative ions from the salt bridge move to the anode to
neutralize the positive charge build-up due to Zn+2 cations
released into solution by the Zn anode.
Features of an Electrochemical Cell
7. Electrons only flow through the external wire (not in the
solutions) while solution ions migrate in the solution.
8. The ideal voltage that can be produced at the start of the
cell’s operation is found by adding the voltages of the half
cell reactions. Note that since Zn is oxidized, the reaction
below must be written in reverse with the sign changed from
-0.76 to +0.76.
Predict Anode, Cathode, Current Direction, Mass
Changes and Initial Voltage
KNO3
solution
Al
Al(NO3)3
solution
Ag
AgNO3
solution
What Is Voltage?
Voltage is a measure of the potential energy carried by
flowing electrons. It is the energy per coulomb (6.24 x 1018)
of electrons. 1 V is 1J/C.
Potential Difference
Electrons do not flow in a single half-cell so an individual
half-cell voltage cannot be determined. However, if two
half cells are connected, the potential difference between
those two cells can be measured.
Iron Half-Cell
Fe (s)
1M
Fe(NO3)3
solution
A Review of Electrical Terminology
A coulomb (C) is a
number of
electrons (1C =
6.24 x 1024 e-).
Electrical current
(amperes – A) is
the number of
electrons passing
per second,
A = #C/s.
Voltage is the energy
carried per
number of
electrons, V = J/C.
Standard Reduction Potentials
All half cell reactions are
assigned voltages in terms of
being connected to a
hydrogen half-cell. A
hydrogen half-cell has
hydrogen gas at 1 atm (101.3
kPa) pressure bubbling into a
glass tube which has an inert
(non-reactive) platinum (Pt)
electrode. There may be a
return tube carrying excess H2
(g) away. The electrode is
suspended in a 1 M solution
of strong acid so that there is
1 M H+ in solution.
Possible Reactions At A “Hydrogen Electrode”
Around the platinum electrode, either a reduction will occur
2H+ (aq) + 2 e -
 H2 (g)
,or the reverse oxidation will occur.
Standard Reduction Potentials
Arbitrarily, the hydrogen half cell is assigned a value of zero
and the other cell’s voltage is determined experimentally
with a voltmeter. The hydrogen half-cell reaction is :
2H+ (aq) + 2 e -
 H2 (g)
; Eo
= 0.000 V
Experimentally Determining a Standard Reduction Potential
Under standard conditions, a second electrode is attached to
a hydrogen electrode with a voltmeter between. The
reading on the voltmeter establishes where on the chart
the half reaction will be placed. Note the placement of the
Ag+ reduction half reaction in the SRP chart.
2Ag+ + H2
 2H+ + 2Ag
; Eo
= 0.80 V
Experimentally Determining a Standard Reduction Potential
In the new reaction shown, H+ is being reduced while Zn is
being oxidized. A voltage of +0.76 V is produced when Zn
is oxidized. To write a reduction potential for Zn2+, the
voltage must be reversed since the half reaction is
reversed. Note the placement of Zn2+ in the SRP chart.
2H+ + Zn 
H2 + Zn2+ 
H2 + Zn2+
2H+ + Zn
; Eo
= 0.76 V
; Eo = -0.76 V
The Meaning of Eo
Eo stands for Standard Reduction Potential where the “ o “
means standard state. The standard state for a half-cell is
25o C, all gases in the cell at 101.3 kPa of pressure, all
elements in the cell in their standard states at 25o C, all
solutions in the cell at a concentration of 1 M.
The Sign of Eo
When a half-reaction is reversed from how it is written in a
table (when an oxidation is occuring), the sign of the Eo is
reversed. The Eo values of the reduction and oxidation
reactions are added to give the voltage value of an
electrochemical cell as it starts under standard conditions.
What The Eo Sign Means
A positive E sign means that
the cell will do work
(produces electricity – the
reactions are spontaneous)
A negative E sign means that
the cell has to be given
energy for the reactions to
occur – the reactions are
not spontaneous.
Al+3 + 3Ag  Al + 3Ag+1 is ?
Zn+2 + Mg  Zn + Mg+2 is ?
Balancing Coefficients Have No Effect on Eo
For the reaction, 3 Ag+ + Al  3 Ag + Al+3,
The half reactions are:
3 Ag+ + 3 e-  3 Ag
;
Eo = +0.80 V
Al  Al+3 + 3 e- ;
Eo = +1.66 V
Eocell = + 2.46 V
Note that the Eo value is unaffected by the coefficient, 3.
The reason for this is that if there are 3 x the electrons capable of
doing 3X the work, the voltage ratio of work/charge remains the same.
3X work / 3X electrons = work / electrons
Area of the Electrodes has No Effect on Eo
The potential difference Eo does not change when the area of
the electrodes is changed.
The reaction rate does increase when the surface area
increases and this releases more electrons per second
which increases the amperage or current flow.
A larger electrode surface area also increases the time that a
cell can operate.
The Size of Eo and Rate of Reaction
In the reaction:
3 Ag+ + 3 e-  3 Ag ;
Eo = +0.80
V
Al  Al+3 + 3 e- ; Eo = +1.66 V
Eocell = + 2.46 V
The size of Eo does NOT indicate
anything about the rate of
reaction. This reaction normally
proceeds very very slowly. This
is because aluminum reacts with
oxygen in a solution to form an
oxide coating (Al2O3) on the
aluminum metal that adheres
strongly and inhibits further
reaction unless sufficient
activation energy is supplied.
Aluminum
oxide
Sil
ver
Aluminum
Detarnishing Silverware
Tarnished silverware is Ag2O which when reacted with Al
causes silver (Ag+) to be reduced to Ag :
3 Ag+ + 3 e-  3 Ag
;
Eo = +0.80 V
Al  Al+3 + 3 e- ;
Eo = +1.66 V
Eocell = + 2.46 V
This reaction is accomplished by heating tarnished
silverware on aluminum foil with baking soda (makes a
weak base that removes the aluminum oxide coat) .
E Values for Nonstandard Conditions
Increasing concentrations of solutions causes a voltage rise
since an increase in reactant concentration tends to shift
equilibrium to the products.
2 Ag+ + 2e-  2Ag ; ½ react. Eo = +0.80
Cu2+ + 2e-  Cu ; ½ react. Eo = -0.34 , Cell Eo = +0.46 V
2 Ag+ +
2e-
Cu2+ + 2e-
 2Ag
 Cu
E > +0.80
E > -0.34
A reduction in reactant concentration causes a shift towards
the reactants ( a greater tendency for the reverse reaction)
which causes a voltage drop from the expected Eo.
2 Ag+ + 2e-  2Ag
Eo < +0.80
Cu2+ + 2e-  Cu
Eo < -0.34
When Cells Reach Equilibrium
When a cell starts, it produces the full theoretical Eo voltage
value. As the cell operates, this voltage gets lower and
lower. The reason for this is that as products build up, the
opposing reactions have a greater tendency. Ex:
2 Ag+ + 2e-  2Ag
Cu2+ + 2e-  Cu
Thus the starting Eo drops for the reduction reaction and also
drops for the oxidation reaction so that the final E is 0.00 V
Selecting Preferred Reactions
In cells where multiple reduction or multiple oxidations are
possible, the reduction and oxidation reaction with the
highest potential (E) will be the ones that occur.
Example 1 of Selecting the Preferred Reaction
Example 2 of Selecting the Preferred Reaction
The Lead-Acid Storage Battery
A charged car battery (Lead-Acid) has electrodes of Pb and
PbO2 . It can produce electric current with the electrodes
getting a lead IV sulfate coating and the reactions can be
reversed by supplying electric current to produce the
original Pb and PbO2 electrodes.
The Reactions in a Lead-Acid Storage Battery
Typical Cell Construction of a 12 V Car Battery
Car batteries typically are made up of six 2 volt cells
connected in series to yield 12 V.
Different Acid Concentration When Discharged
As a lead-acid storage battery discharges, sulfuric acid is
used up and water is produced. This reduces the
solution’s acid concentration and its density becomes less
(from 1.30 to 1.10 g/mL). A floating hydrometer will float
lower in a discharged battery.
Determining Which Electrode is the Anode/Cathode
The anode is the electrode where oxidation is occurring while
the cathode is where reduction is occurring. The
menumonic, “An Ox CaRed” can be used to remember this.
The anode in an electrochemical cell is negative while the
cathode is positive.
The Zinc-Carbon Battery
Fuel Cells
The electron flow shown in this diagram is wrong and should
be in the opposite direction.
Ballard Hydrogen Oxygen Fuel Cell
Current (electrons) flows from the anode (-), where hydrogen
gas is releasing electrons (oxidized) to become positive
protons, to the cathode (+) where oxygen is accepting
electrons (from the current) and protons to make water.
Corrosion of Iron (Called Rusting)
Deep in the centre of a water drop, which is oxygen-poor, iron
metal oxidizes to ions, releasing electrons (Fe  Fe2+ + 2e-).
The electrons migrate to the outer edges where they reduce
dissolved oxygen (in the oxygen-rich outer region of the
drop) to hydroxide ions (1/2 O2 + H2O +2e -  2OH-). The
2OH- ions react with the Fe2+ to produce insoluble Fe(OH)2 at
the outer edges of the water drop. Fe(OH)2 is then oxidized
to Fe2O3 . XH2O (rust) due to further contact with oxygen.
½ O2 + H2O + 2e-  2OH-
Fe2+ + 2OH-  Fe(OH)2
Fe  Fe2+ + 2e-
Preventing Corrosion
Generally, three things can be used to prevent metals from
corroding.
1. Use a metal whose oxide layer adheres strongly to the
metal. The oxide coating then prevents water and oxygen
from making contact with the metal. Aluminum and copper
produce oxides that bind to the metal and protect it from
further oxidation.
Preventing Corrosion
2. Another method of preventing corrosion is to apply a
protective coating or a paint coating to a metal which then
prevents oxygen and water from reacting with the metal.
Preventing Corrosion
3. Another method of preventing corrosion is to plate a second
metal onto a metal that tends to corrode. The second, plated
metal tends to oxidize easier but has an oxide that strongly
adheres to the second metal forming a barrier to oxygen and
water. Steel can be plated with tin whose oxide (SnO2) forms a
protective layer. Galvanized nails are iron coated with zinc
whose oxide (ZnO) also strongly adheres as a protective layer.
Cathodic Protection
Cathodic protection means that a metal to be protected against
corrosion (like iron) is connected to another metal (like
magnesium) that has a greater tendency to oxidize. Since the
other metal (magnesium) is oxidized before the protected
metal (iron) it is connected to, the protected metal receives
electrons (acting as a cathode) from the metal giving
electrons (the anode) [An Ox Cared]. In essence the
magnesium is sacrificed to prevent iron from oxidizing.
Cathodic Protection
Some pipes, ships’ hulls and vehicles are cathodic protected
by putting a low voltage (low energy) current into them,
making the metal unable to give off electrons (corrode). The
current provides electrons to surrounding water and oxygen
rather than the metal.
Cathodic Protection
Sacrificial anodes and impressed electrical current is
commonly used to protect underground metal storage
tanks.
Changing Chemical Surroundings to Fight Corrosion
The reduction reaction of oxygen causes the oxidation of
metals like iron.
½ O2 + 2H+ (10-7 M) + 2e-  H2O ; E= 0.82 V
If oxygen can be removed from a solution in contact with a
metal, the only reduction that can occur higher than iron is
2 H+ (10 -7 M) + 2e-  H2 (g) ; E = -0.45 V which is a much
slower reaction.
Even better protection is obtained by adding OH- ions (NaOH)
to a solution around the metal. The reduction half reaction
is now 2 H2O + 2e-  H2 (g) + 2 OH- ; E = -0.83 V
which is below iron in the SRP chart.
End of Presentation