Electrolytic Cells
In electrolysis we use electrical energy to bring about chemical change. An otherwise non-spontaneous reaction is “driven” in the forward
direction by application of a voltage.
Electrolytic Cell: Electrical energy from an external source drives a non-spontaneous reaction.
Examples of electrolysis:
1. Recharging a secondary battery.
2. Electrolysis of molten salts.
3. The breakdown of water into O2 and H2 gases.
4. Electrolysis of aqueous salts.
5. Electroplating of metals.
6. Purification of metals.
Voltaic Cell-Discharge
The processes occurring during the
discharge and recharge of a lead-acid
battery.
Electrolytic Cell-Recharge
As with a voltaic cell, oxidation occurs at the anode and
reduction takes place at the cathode. An external voltage
source supplies the cathode with electrons, which is
negative, and removes them from the anode, which is
positive.
Electrochemistry
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Electrolysis of Molten salts
Molten salts can be transformed into their elements by electrolysis, a process that, splits the ionic compound
decomposing it to its elements. Inert electrodes are used in these cases.
Conditions are non-standard! Standard Reduction Potentials do not apply! Not an aqueous system!
Voltage Source
Example: Electrolysis of molten NaCl.
cathode
What half reaction occurs:
at the anode?
at the cathode?
anode
+
–
+
–
What is the overall reaction with phase
labels? (The melting point of NaCl is
801°C, the melting point of Na is 97.8°C.)
Note the sign conventions
of anode and cathode.
Question:
A mixture of KI and LiBr is melted and electrolyzed. What products are formed at the lowest applied voltage that will
produce a reaction? What is the overall cell reaction? During electrolysis of a mixture of molten salts
- the more easily oxidized species (stronger reducing agent) reacts at the anode.
How do you decide?
- the more easily reduced species (stronger oxidizing agent) reacts at the cathode.
How do you decide?
Electrochemistry
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Electrolysis of Water
Electrolysis of water decomposes (splits the water) into its
elements.
Anode
Cathode
2 H2O(l) —> 2 H2(g) + O2(g)
A 9 V battery is sufficient.
An inert salt (one that will not react) must be present in the water
to transport charge.
Oxidation takes place at the anode:
2 H2O(l) —> O2(g) + 4 H+(aq) + 4 e–
Reduction takes place at the cathode:
2 H2O(l) + 2 e– —> H2(g) + 2 OH–(aq) a
In practice, the half-reaction voltages required are as follows
and are not the same as in a table of SRP’s:
(-) Reduction:
2 H2O(l) + 2 e– —> H2(g) + 2 OH–(aq)
Ered (Actual “Over voltage”) ≈ –1 V
(E°red -0.83 V)
(+) Oxidation:
2 H2O(l) —> O2(g) + 4H+(aq) + 4 e– Eox (Actual “Over voltage”) ≈ –1.4 V
(E°ox -1.23 V)
These are the voltages we will use when analyzing a aqueous system.
What is the “actual-voltage” that must be applied (Ecell) for the electrolysis of water?
Electrochemistry
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Electrolysis of AQUEOUS Salts
Under aqueous conditions we must include the possible oxidation or reduction of water as well as
the salt ions. When an aqueous salt solution is electrolyzed:
- The strongest oxidizing agent (most easily reduced) is reduced, and
Anode
- The strongest reducing agent (most easily oxidized) is oxidized.
For example consider the electrolysis of
1 M NaI(aq):
Possible cathode reactions:
a) Na+ + e- —> Na(s)
E°red = -2.71 V
Wire mesh
b) 2 H2O(l) + 2 e- —> H2(g) + 2 OH–
Ered ≈ -1 V (nonstandard)
(Zoom View)
Possible anode reactions:
a) 2 I– —> I2(s) + 2 eE°ox = -0.54 V
–
+
1M NaI(aq)
b) 2 H2O(l) —> O2(g) + 2 H+ + 2 eEox ≈ -1.4 V (nonstandard)
Determine the net redox reaction at the
minimum voltage needed to produce a
reaction.
Electrodes are inert. Therefore, there is
no oxidation of the anode electrode.
Electrochemistry
What voltage is needed?
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Electrolysis of AQUEOUS Salts
Determine the products that will form, using the minimum voltage needed to produce a
reaction, for electrolysis of 1 M SnCl2(aq) at 25°C,
+
(-) All Possible Cathode Rxns:
1 M SnCl2(aq)
(+) All Possible Anode Rxns:
Electrodes are inert. Therefore, there is no
oxidation of the anode electrode.
What is the net reaction and expected minimum voltage needed:
Electrochemistry
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Electrolysis of AQUEOUS Salts
Write the balanced redox reaction that occurs at the minimum voltage needed to produce a reaction
when the electrolysis of 1 M Cu(NO3)2(aq) is conducted using:
a) A Pt anode and a Pt cathode
b) A Cu anode and Cu cathode
Electrochemistry
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Overview of Electrical Units
Energy (J): 1 J = kg•m2/s2
Coulomb (C):
fundamental unit of electric charge.
One electron has a charge of 1.602x10-19 C.
1 C = 6.241x1018 e– = 1.0364x10-5 mole e–
Voltage (V):
the emf the electrons “feel”.
1 V = 1 J/C
Faraday (F):
absolute charge of one mole of electrons.
F = 9.6485x104 C/mole
Current (A): the number of electrons that “flow” per second.
ampere (A or Amp), 1 A = 1 C/s
Work (w):
Voltaic cell (system can do work):
The maximum work that can be done: wmax = -nFEcell = ∆G ( units are J)
Electrolytic cell (work done on system by Eexternal): w = nFEext ( units are J)
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Stoichiometry of Electrolysis
Faraday’s Law of Electrolysis states that the amount of substance produced at each electrode
is directly proportional to the quantity of charge flowing through the cell. (This assumes 100%
efficiency, not achievable.)
AMOUNT (mol)
of substance
oxidized or
reduced
(or mol product)
balanced
half-reaction
AMOUNT (mol)
of electrons
transferred
Faraday constant
(9.6485x 104 C/mol e−)
M (g/mol)
MASS (g)
of substance
oxidized or
reduced
CHARGE
(C)
Practice Problem: How many minutes does it take to form 10.0 L of
oxygen gas at 28°C and a total pressure of 1.00 atm by electrolysis of
water if a current of 1.3 A is passed through the electrolytic cell. Assume
100% efficiency. The vapor pressure of water at 28°C is 28.3 torr.
time (s)
Recall:
1 Amp = 1 C/sec
CURRENT
(A)
(or g product)
What mass of hydrogen gas forms?
Electrochemistry
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Electrolytic Refining (Purification) of Metals
Example: Refining of Copper
In the electrolytic refining of copper a current of 1.50 A is passed through the cell for
2.00 hours. If the cell operates at 75.0% efficiency, what mass of pure copper is
obtained?
Electrochemistry
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Electroplating - Surface Coating of a Metal
Ni(s) is oxidized at the anode.
Ni2+(aq) is reduced at the cathode.
Practice Problem: A piece of steel is to be electroplated with nickel by
the electrolysis of a NiSO4(aq) solution.
a) If the process is 80.0% efficient, what mass of Ni is plated out if a
current of 374 mA is maintained for 3.80 hours?
Net Reaction?
b) If the cathode has a surface area of 100 cm2, what average thickness
of Ni metal was deposited? The density of Ni is 8.91 g/cm3.
c) How much work was done on the system if the external voltage was
5.76 V?
During electrolysis, Ni atoms
are transferred from the nickel
anode to the steel cathode,
plating the steel with a thin
layer of nickel atoms.
E°cell = 0 V
Only a small emf is needed!
d) Is this amount of work done on the system equal in magnitude to the
total energy lost by the voltage source? Explain your answer.
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Electrical Energy and Electrolysis
Practice Problem:
a) Calculate the total energy in Joules required to produce 1.00 kg of Mg from the electrolysis of molten MgCl2 if the
applied emf is 5.00 V. Assume that the process is 90% efficient.
b) If the time to complete the process was 52.0 minutes, what was the average current flow during this time?
c) Can Mg be produced by electrolysis of a 1 M MgCl2(aq) solution? Provide clear reasoning for your answer.
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