Electrolytic Cells use an external power supply to force a non-spontaneous redox
reaction to occur.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn(s)
E   2.12 V
To get an idea of what an electrolytic cell, in general is, we’ll start by looking at two
half-reactions and the overall redox reaction we get by adding them.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn(s)
E   2.12 V
We’ll start with the half-reaction we get by reversing the reduction of chlorine on
the table, to get the (click) oxidation of chloride ions
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn(s)
E   2.12 V
Notice, because we reversed the equation, (click) the sign on the E naught is
switched from positive to negative.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s )
E   0.76 V
Cl 2( g )  Zn(s )
E   2.12 V
Next, we’ll add the half-reaction for the reduction of zinc ions, the way it is on the
table, with an E naught value of negative 0.76 Volts.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g )  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn(s)
E   2.12 V
We’ll add these two half-reactions to get the overall redox equation.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g )  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn(s)
E   2.12 V
electrons gained = electrons lost
Notice the electrons gained by the zinc ion, are (click) equal to the electrons lost by
the chloride ions.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g )  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn(s)
E   2.12 V
Therefore, electrons can be cancelled before adding the half-reactions.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Cl 2(g)  Zn( s)
E   2.12 V
On the left side we have Zn2+ plus 2 Cl minus
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn( s)
E   0.76 V
Zn(s)  Cl 2( g)
E   2.12 V
And On the right side we have Zn solid plus Cl2 gas
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Zn(s)  Cl 2(g)
E   2.12 V
To get the E naught value for the overall redox equation, (click) we add –1.36 volts
and –0.76 volts, to give us (click) –2.12 volts
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Zn(s)  Cl 2(g)
E   2.12 V
This redox reaction is
non-spontaneous
A negative E naught value for the overall redox reaction, means that (click) this
redox reaction is non-spontaneous as written.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g)  2e

E   1.36 V
Zn(s)
E   0.76 V
Zn(s)  Cl 2(g)
E   2.12 V
This redox reaction is
non-spontaneous
So that means, if we simply mixed Zn2+ and Cl–, we would (click) NOT get solid zinc
and chlorine gas.
Consider the following half-reactions
and overall redox reaction:
2Cl

Zn 2  2e 
Zn
2
 2Cl

Cl 2(g )  2e

E   1.36 V
Zn(s)
E   0.76 V
Zn( s)  Cl 2(g)
E   2.12 V
But we can Force this reaction to occur, using a process called electrolysis.
Electrolysis uses an external
power supply or battery to force
a non-spontaneous redox
reaction to occur.
Zn
2
 2Cl

Zn(s)  Cl 2(g)
E   2.12 V
Electrolysis uses an external power supply or battery to force a (click) nonspontaneous redox reaction to occur.
Inert
Carbon
Electrode
Inert
Carbon
Electrode
Electrolysis takes place in an electrolytic cell. We’ll look at the simplest type of electrolytic cell. This is what we
call a Type 1 Electrolytic cell. We’ll start off with (click) two inert carbon electrodes in a single container.
Power
Supply
Inert
Carbon
Electrode
+ –
Then we add a direct current power supply and wires.
Inert
Carbon
Electrode
Power
Supply
+ –
Neutral
+
e–
e–
+
e–
+
e–
+
e–
+
+
+
+
e–
e–
e–
+
e–
+
e–
+
e–
+
e–
Neutral
e–
+
e–
+
e–
+
e–
+
Before the power supply is connected, the electrodes are both neutral. Using simplified symbols, we’ll
represent a few protons by + signs and a few electrons as “e”s with a negative charge. Because they are neutral,
Power
Supply
+ –
Positive
Neutral
+
e–
e–
+
e–
+
e–
+
e–
+
+
+
+
e–
e–
e–
+
e–
+
e–
+
e–
+
e–
Negative
Neutral
e–
+
e–
+
e–
+
e–
+
When we close the switch and connect the power supply, it takes electrons from the electrode attached
to the positive terminal and pumps them on to the electrode attached to the negative terminal
Power
Supply
+ –
Positive
Has a deficiency
of electrons
+
+
+
+
+
e–
e–
+
e–
+
+
e–
+
e–
e– –+
+ e e–
e– –+
+ e e–
e– –+
+ e e–
e–e–+
The electrode on the left now has a deficiency of electrons, giving it a net positive
charge
Power
Supply
+ –
Positive
Has a deficiency
of electrons
+
+
+
+
+
e–
e–
+
e–
+
+
e–
+
Negative
e–
e– –+
+ e e–
e– –+
+ e e–
e– –+
+ e e–
e–e–+
Has an excess of
electrons
And the electrode on the left now has an excess of electrons, giving it a net
negative charge
Power
Supply
Positive
+ –
+
ANODE
+
+
+
+
e–
e–
+
e–
+
+
e–
+
e–
e– –+
+ e e–
e– –+
+ e e–
e– –+
+ e e–
e–e–+
In an electrolytic cell, the electrode attached to the + terminal of the power supply
is called (click) the Anode
Power
Supply
Positive
+ –
+
ANODE
Negative
+
+
+
+
e–
e–
+
e–
+
+
e–
+
e–
e– –+
+ e e–
e– –+
+ e e–
e– –+
+ e e–
e–e–+
CATHODE
And the electrode attached to the negative terminal of the power supply is called
(click) the Cathode.
Power
Supply
Positive
+ –
+
ANODE
A+
Negative
+
+
+
+
e–
e–
+
e–
+
+
e–
Just remember A plus for anode,
+
e–
e– –+
+ e e–
e– –+
+ e e–
e– –+
+ e e–
e–e–+
CATHODE
Power
Supply
Positive
+ –
+
ANODE
A+
Negative
+
+
+
+
e–
e–
+
e–
+
+
e–
+
e–
e– –+
+ e e–
e– –+
+ e e–
e– –+
+ e e–
e–e–+
CATHODE
C–
And C minus for Cathode. Remember, this ONLY works for an Electrolytic cell, not
an electrochemical cell.
Power
Supply
+ –
ANODE
+
e–
e–
+
e–
e–
e–
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
–
CATHODE
–
We’ll redistribute the electrons and stop showing the protons, for simplicity.
Power
Supply
+ –
ANODE
+
e–
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
e–
+
–
e–
e–
e–
ZnCl2(l)
We’ll add some molten Zinc chloride to the container
CATHODE
–
Power
Supply
+ –
ANODE
+
e–
e–
+
e–
Cl–
e–
e–
Cl–
Zn2+
Cl–
Cl–
Zn2+
Cl–
Zn2+
Cl–
Cl–
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
–
CATHODE
–
Cl–
Zn2+
ZnCl2(l)
A molten salt like zinc chloride consists of ions that are in constant random
motion.
Power
Supply
+ –
ANODE
+
e–
e–
+
e–
e–
Cl–
e–
Zn2+
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
–
Cl–
ZnCl2(l)
Now we’ll focus on one zinc ion and two chloride ions.
CATHODE
–
Power
Supply
+ –
ANODE
+
e–
e–
+
e–
e–
Cl–
e–
Zn2+
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
–
CATHODE
–
Cl–
ZnCl2(l)
The positive zinc ions will be attracted to the negative cathode while the negative
chloride ions will be attracted to the positive anode.
Power
Supply
+ –
ANODE
+
e–
e–
+
e–
Cl–
e–
e–
Zn2+
Cl–
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
ZnCl2(l)
Now, we’ll concentrate on the cathode and see what happens there.
–
Power
Supply
+ –e
ANODE
+
––
e–
e–
+
e–
Cl–
e–
e–
Zn2+
Cl–
ZnCl2(l)
The zinc ion will gain 2 electrons from the cathode,
e– –
e –
e
e– –
e –
e
e– –
e –
e
–
e e–
–
CATHODE
–
Power
Supply
+ –
ANODE
+
e– –
e
e–
+
e–
Cl–
e–
e–
e–
Zn2+
–
Zn
e
Cl–
ZnCl2(l)
And turn into a zinc atom
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
e– –
e
e–
+
e–
Cl–
Zn
e–
e–
Cl–
ZnCl2(l)
The equation for what just happened is Zn2+,
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
e– –
e
e–
+
e–
Cl–
e–
e–
Zn
Cl–
ZnCl2(l)
Plus 2 electrons
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
e– –
e
e–
+
e–
Cl–
e–
e–
Zn
Cl–
ZnCl2(l)
Gives Zinc solid
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
Reduction
e– –
e
e–
+
e–
Cl–
e–
e–
Zn
Cl–
ZnCl2(l)
This is an example of reduction.
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
Reduction
e– –
e
e–
+
e–
Cl–
e–
e–
Zn
Cl–
e–
e–
CATHODE
–
e– –
e –
e
e– –
e –
e
–
e e–
ZnCl2(l)
So we see that reduction of the cation Zn2+ occurs at the cathode
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
Reduction
e– –
e
e–
+
e–
Cl–
e–
e–
Zn
Cl–
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
ZnCl2(l)
Now we’ll have a look at the anode and see what happens.
CATHODE
–
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ee––
ANODE
+
e–
+
e–
e–
e–
Reduction
e– –
e
e–
Cl
Cl–
Zn
e–
Cl–
Cl
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
ZnCl2(l)
Each chloride ion will lose one electron and change into a chlorine atom (Click half
way through)
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
ANODE
+
Reduction
e– –
e
e–
+
e–
e–
e–
Cl
Zn
e–
Cl
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
ZnCl2(l)
The 2 chlorine atoms then join to form a molecule of Cl2 or chlorine gas
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
Reduction
e– –
e
e–
+
e–
e–
e–
Cl
Zn
e–
Cl
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
ZnCl2(l)
The process taking place on the anode can be summarized by the equation (click),
2 Cl minus
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
Reduction
e– –
e
e–
+
e–
e–
e–
Cl
Zn
e–
Cl
ZnCl2(l)
Gives Cl2 gas
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
Reduction
e– –
e
e–
+
e–
e–
e–
Cl
Zn
e–
Cl
ZnCl2(l)
Plus 2 electrons.
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
Reduction
e– –
e
e–
+
e–
e–
e–
Cl
Zn
e–
Cl
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
ZnCl2(l)
The process of chloride ions losing electrons (click) is called oxidation
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
Reduction
e– –
e
e–
+
e–
e–
e–
Cl
Zn
e–
Cl
e–
e–
–
e– –
e –
e
e– –
e –
e
–
e e–
CATHODE
–
ZnCl2(l)
And we see that oxidation of chloride ions to chlorine gas takes place at the anode.
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
+
Reduction
CATHODE
Zn2+
Zn
–
ZnCl2(l)
So again zinc ions are reduced at the cathode to form zinc atoms.
–
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
+
Zn2+ + 2e–  Zn(s)
+
Reduction
CATHODE
Cl–
Cl
Zn
Cl
Cl–
–
ZnCl2(l)
And chloride ions are oxidized at the anode to form chlorine gas, Cl2.
–
Power
Supply
2Cl–  Cl2(g) + 2e–
+ –
Oxidation
ANODE
Zn2+ + 2e–  Zn(s)
Reduction
e–
CATHODE
+
Zn
–
2Cl–  Cl2(g) + 2e–
–
Zn2+ + 2e–  Zn(s)
ZnCl2(l)
So as this cell operates, we can visualize zinc metal growing on the surface of the
cathode and bubbles of chlorine gas forming on the anode
2Cl–  Cl2(g) + 2e–
Oxidation
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
Reduction
CATHODE
ANODE
+
Zn
–
2Cl–  Cl2(g) + 2e–
Product at Anode
Zn2+ + 2e–  Zn(s)
ZnCl2(l)
So we can say the product at the anode is chlorine gas
–
2Cl–  Cl2(g) + 2e–
Oxidation
Power
Supply
Zn2+ + 2e–  Zn(s)
+ –
Reduction
CATHODE
ANODE
+
Zn
–
2Cl–  Cl2(g) + 2e–
Product at Anode
Zn2+ + 2e–  Zn(s)
ZnCl2(l)
And the product at the cathode is zinc metal.
Product at Cathode
–
Molten zinc chloride is electrolyzed.
ZnCl 2( l ) 
2
Zn( l )
Reduction
at Cathode


2Cl ( l )
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
Now that we’ve seen how this cell works, we’ll show you a process you can use for any questions
involving the electrolysis of a molten salt. We’ll use our example here of molten zinc chloride.
Molten zinc chloride is electrolyzed.
ZnCl 2( l ) 
2
Zn( l )
Reduction
at Cathode


2Cl ( l )
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
Molten salts always consist of mobile positive and negative ions, so we write the
dissociation equation for the salt forming liquid ions. We have ZnCl2 liquid
Molten zinc chloride is electrolyzed.
ZnCl 2( l ) 
2
Zn( l )
Reduction
at Cathode


2Cl ( l )
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Forms Zn2+ liquid
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
Molten zinc chloride is electrolyzed.
ZnCl 2( l ) 
2
Zn( l )
Reduction
at Cathode


2Cl ( l )
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Plus 2 Cl minus liquid
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
We write C minus for cathode, underneath the positive ion. The positive ions are
attracted to the negative cathode.
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
And we write A+ for anode, underneath the negative ions. The negative ions are
attracted to the positive anode.
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Reduction of Zn2+ ions occurs at the cathode.
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Its half-reaction is: Zn2+ + 2e–  Zn(s) .
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
And the product at the cathode is zinc solid
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Oxidation of Cl– ions occurs at the anode
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Its half-reaction is: 2Cl– Cl2(g) + 2e– .
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
And the product at the anode is chlorine gas
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
We get the equation for the overall redox reaction by adding up the two halfreactions. It is (click) Zn2+ + 2Cl–  Zn(s) + Cl2(g)
ZnCl 2( l ) 
2
Zn( l )


2Cl ( l )
C–
A+
Reduction
at Cathode
Oxidation
at Anode
Zn 2  2e   Zn(s) 2Cl   Cl 2(g)  2e 
Product at Cathode is Zn(s)
Overall Redox Reaction:
Product at Anode is Cl2(g)
Zn 2  2Cl   Zn(s)  Cl 2(g)
Going through this process should greatly help you with any questions you get
involving the electrolysis of molten salts, or what we call Type 1 electrolytic cells.