9.5/19.2 ELECTROLYTIC CELLS and ELECTROLYSIS
Electrochemical Cells
Spontaneous
(Galvanic Cells
aka Voltaic Cells)
Non-Spontaneous
(Electrolytic Cells)
*ELECTROLYTIC CELLS:
 make non-spontaneous redox reactions occur by providing energy in the form of electricity from an external source
 involve the passing of electricity through an electrolyte (Recall: An electrolyte is a substance that does not conduct when solid, but
does conduct when molten or in aqueous solution.)
 involve the conversion of electrical energy into chemical energy
*ELECTROLYTIC CELL COMPONENTS:
i) two __________________;
ii) an ___________________;
iii) an external ______________or power source (power source acts as an “electron pump”,(i.e power source _________ an electron
transfer)
OVERVIEW: VOLTAIC vs. ELECTROLYTIC Cells
VOLTAIC Cell
CHARACTERISTIC
ELECTROLYTIC Cell
Spontaneity
Standard Cell potential (ΔE˚)
ΔE =
ΔE =
Energy Conversion
Chemical energy  electrical energy
(i.e. chemical reaction produces and
electric current)
 Cathode: ______________ terminal
 Anode: _______________ terminal
 Cathode: ________________
 Anode: _________________
Electrons flow from ____________ to
________________.
(*from -ve to +ve terminal)
Electrical energy  chemical energy
(i.e. electrical current produces a
chemical reaction)
 Cathode: ______________ terminal
 Anode: _______________ terminal
 Cathode: ______________
 Anode: _______________
Electrons flow from ____________ to
________________.
(*from +ve to –ve terminal)
Cations to cathode (via salt bridge)
Anions to anode(via salt bridge)
Cations to cathode (-ve terminal)
Anions to anode(+ve terminal)
*LABELLING
(Cathode VS Anode)
HALF RXNS
(Cathode VS Anode)
Direction of e- movement
Salt Bridge?
Direction of ion movement
*What are the main differences between the two types of cells? What is the same?
*ELECTROLYSIS
 the process of supplying _________________ energy to force a ______-spontaneous reaction to occur
*Applications:
i) Production of Elements (e.g. Na, Li, Al, Cl2)
ii) Refining of metals
iii) Electroplating
A. ELECTROLYSIS OF MOLTEN SALTS
Recall: An electrolytic cell uses ______________________ to force (an otherwise non-spontaneous)
reaction to occur. This is opposite to the process occurring in __________________ cells.
When a molten salt is electrolyzed:
 the positive ______________ migrates toward the negative_________________ where it is ________________;
 the negative anion migrates toward the _______________ anode where it is _________________.
Because there are not other elements around, and salts have a metal and non-metal, it’s easy to figure out the product.
e.g. sodium chloride melts at 801° C as follows:
NaCl(l) 
Cathode (-) half rxn:
Anode (+) half rxn:
Overall redox rxn:
*always use state symbols
or
*This represents one of the many important applications of _____________________ – the production of ___________________..
B. ELECTROLYSIS OF AQUEOUS SOLUTIONS
*When a solution is electrolyzed, we must now consider the presence of water. Water can act as a __________________agent or an
_____________________ agent. It’s time to use the two water reactions you highlighted previously (at -0.83 V and +1.23 V).
Remember the following:
**Water is more easily reduced (i.e. it is a better oxidizing agent) at -0.83 V than anything above it. (i.e. If your solution contains
cations with a more negative value (above in the table), then you get H2(g) instead from the reduction of water.)
and
**If the solution contains anions with a more positive value than +1.23 V (below in the table), when you obtain O2(g) from the
oxidation of water.
(IB Chemistry considers that most complex ions (SO42-, NO3- etc.) are stable towards oxidation (i.e. not readily oxidized), and
therefore water is oxidized instead when these ions are present.
Special Exception: CHLORIDES
*A dilute solution of chloride ions will generate oxygen at the anode (as expected from the rules above). Usually we deal with
standard (1.0 M) solutions. However, when dealing with concentrated solutions we can expect some variation. **A concentrated
solution of NaCl(aq), (called “brine”) will yield chlorine gas, Cl2(g) instead.
Demo A: ELECROLYSIS OF WATER (The SOA still gets reduced at the cathode!)
a) Why must an electrolyte such as Na2SO4 be added?
b) Write the half reaction occurring at:
i) the cathode(-ve):
ii) the anode(+ve):
c) Add the half reactions:
Demo B: ELECTROLYSIS OF BRINE (i.e. salt, NaCl, water)
a) What is different? What do you smell forming at the anode?
b) Write the half reaction occurring at:
i) the cathode(-ve):
ii) the anode(+ve):
c) Add the half reactions:
*Practice: ELECTROLYSIS:
Determine the products at each electrode for the following electrolytes.* Pay close attention to the state symbols!
Electrolyte
NaCl(l)
Cathode(-ve terminal) Product
Anode(+ve terminal) Product
NaCl(aq) dilute
NaCl(aq) concentrated
CuSO4(aq)
Na2SO4(aq)
MgBr2(aq)
PbBr2(l)
PbBr2(aq)
HINTS: Summary of How to Quickly Obtain ∆E°cell and the Overall Redox Rxn
A) For Spontaneous Voltaic Cells:
i) Obtaining ∆E°cell
Recall: ∆E°cell = E°r (cathode) - E°r (anode) = ______________________ value
Therefore: subtract lower E°r from higher E°r . [i.e. E°r (higher) - E°r (lower)]
ii) Obtaining overall redox reaction
Reverse the half reaction with the __________________ E°r (i.e. this represent the reaction taking place at the
____________________/ i.e.the oxidation rxn) and add it the other unchanged half reaction (i.e. the cathode/reduction
rxn).
*Make sure: number of e- gained = number of e- lost!
B) For Nonspontaneous Electrolytic Cells:
i) Obtaining ∆E°cell
Recall: ∆E°cell = E°r (cathode) - E°r (anode) = ___________________value
Therefore: subtract higher E°r from lower E°r . [i.e. E°r (lower) - E°r (higher)]
ii) Obtaining overall redox reaction
Reverse the half reaction with the _________________ E°r (i.e. the anode/oxidation rxn) and add it the other unchanged
half reaction (i.e. the cathode/reduction rxn). *Again, make sure: number of e- gained = number of e- lost.
FARADAY’S LAW (Section 10.7 also)
 the amount of a substance produced or consumed in an electrolytic reaction is __________ proportional to the quantity of
_____________ that flows through the circuit
*Determining the quantity of electricity charge:
q=
Where: q = quantity of ___________ in ____________;
I = electric ___________ in _____________ (1 A = 1 ______)
t = time in __________
*Faraday’s Constant (F)
 the charge of one _______ of __________ (F = ____________________)
*Determining the number of moles of electrons
:ne-=
q
F
;
ne-=
It
F
Half-Cell Stoichiometry
Examples:
1. What is the mass of Aluminum deposited at the cathode of a aluminum electrorefining cell operating at 12.0 A for 40.0 min?
Steps:
i) Write balanced half-reaction.
ii) Convert given measurements to moles
iii) Calculate the amount of the required substance using the mole ratio
iv) Convert the calculated quantity to the quantity required (i.e. answer the question)
2. In a copper electroplating cell, 0.275 g of copper is to be deposited from a copper(II) sulfate solution in a time of 20.0 min. Predict
the current required.
Factors Affecting Amount of Electrolysis Product:
Factor
Increase current
Result
more product
Increase time
more product
*Charge on ion
the greater the charge, the more electrons per reduction, therefore the less
product formed (see example below)
*Try the following:
Three moles of electrons are supplied to each of the following ions: Na+, Mg2+ and Al3+. How many moles of each
product will be formed? Write out reduction half reactions for each.