Electrochemistry
Chapter 18
Electrochemistry
• Electrochemistry
– the branch of chemistry that studies the electricityrelated application of oxidation-reduction reactions
– Redox reactions involve the transfer of electrons
from the oxidized substance to the reduced
substance
This picture shows the lab
we did – we placed zinc into
CuSO4 and the Cu+2 came out
of solution (was reduced) and
deposited copper metal (the
black substance) onto the zinc
metal. The zinc metal was
oxidized and went into solution.
Electrochemistry
• If we had taken the temperature of the CuSO4/Zinc
solution during the reaction, we would have found
that the temperature increased.
• However, if we separate the two metals by a porous
barrier and connect them by a wire, we can change
that heat energy into electrical energy.
This is called a
“Daniel” cell after its
inventor, John
Frederic Daniell, a
British chemist
Electrochemistry
• Electrode: metal conductor used to establish electrical
contact with the non-metallic (solution) part of the circuit.
• Electrolyte: a substance that dissolved in water to give a
solution that will conduct electricity (positive and negative
ions)
Half-cell: a single
electrode in a solution of
its own ions. One halfcell contains the
oxidation reaction and
one contains the
reduction reaction.
Electrochemistry
• Anode: the electrode where oxidation
takes place
• Cathode: the electrode where reduction
takes place
RED CAT and AN OX
Electrochemistry
• Electrochemical Cell
– A system of electrodes and electrolytes in which either
chemical reaction produces electrical energy by the
flow of electrons
In the cell on the right:
What are the electrodes made
of?____________________________
_______________________________
_______________________________
What are the
electrolytes?____________________
_______________________________
____________________________
Electrochemistry
• There are two kinds of electrochemical
cells:
1. Voltaic
2. Electrolytic
Voltaic Cells
• Uses redox reactions that are
spontaneous (they happen by
themselves without any outside energy
input)
The electrolyte
solutions used are
Copper Sulfate and
Zinc Sulfate
Voltaic Cells
• How do we know if the reactions will be
spontaneous? We check Table J (Activity
Series) to see if one metal can displace the
other.
According to Table J, can
zinc displace copper in
the cell to the right?
Does this make the
reaction spontaneous?
Voltaic Cells
Remember: RED CAT and AN OX
LEO the lion goes GER
In this cell, cations (+) in solution are
reduced (gain electrons) at the surface of
the cathode (RED CAT) to become metal
ions.
In this cell, the reduction half
reaction is:
Cu+2(aq) + 2e- → Cu(s)
Voltaic Cell
RED CAT and AN OX / LEO the lion goes GER
Electrons given up by the zinc at the anode
travel through the wire to the cathode.
However, without the salt bridge, the circuit
would not be complete and would not work.
Voltaic Cell
Anions (negatively charged ions) move toward the
anode to replace the electrons that are flowing
through the wire.
Cations move toward the cathode since the positive
charges have been lost during reduction.
Voltaic Cell
• This occurs as the sulfate ions (SO4-)
move through the salt bridge to the anode
and the Na+2 ions move through to the
cathode.
The solution in
the salt bridge
is sodium
sulfate
Voltaic Cells
Now let’s look at it all together:
1. Zinc loses two electrons at the anode and forms Zn+2
(which goes into solution)
2. The electrons flow through the wire to the copper cathode
where it gains two electrons to form solid copper metal.
3. Ions flow through the salt bridge to replenish the lost
charges.
4. The flow of electrons through the wire generates
electricity.
Electrolytic Cells
If, according to Table J, the reaction is NOT
spontaneous, we can force the reaction to
happen by applying a power source, such
as a battery, in the wire between half cells.
Electrolytic Cells
• The process in which an electric current is used to
produce a redox reaction is called electrolysis.
• In electrolysis, the cathode and the anode SWITCH.
Electrolytic Cells
In this cell:
1. Electrons from the negative end of the battery flow through
the wire to the electrode, which is NOW NEGATIVE and
BECOMES THE CATHODE.
2. The electrode attached to the positive end of the battery
loses electrons to it and it is NOW POSITIVE and
BECOMES THE ANODE.
Electrolytic Cells
Let’s put it all together:
1.
Electrons move from the battery to the cathode.
2.
Positive ions (potassium ions) move to the cathode to gain electrons and
become reduced (GER).
3.
Negative ions (chlorine ions) move toward the anode and give up electrons to
replace the ones lost to the battery. They become oxidized (LEO).
4.
The electrons go through the wire back to the battery (positive end).
K+ + 1 e-→ K
RED CAT
at Cathode
2Cl- → Cl2 + 2eAN OX
at Anode
K+ Cl-
Electroplating
An electrolytic cell is used for
electroplating.
- an electrolytic process by which a metal
ion is reduced and solid metal is
deposited on a surface
Electroplating
To plate a silver spoon: make the spoon the cathode (GER).
The silver ions from anode will gain an electron at the
cathode and silver metal will deposit on the spoon.
The electrolyte is Silver Nitrate which provides
more silver ions (Ag+).
+
AgNO3
solution
Comparison of Cells
• Both have:
– Anode – where oxidation
takes place (AN OX)
– Cathode – where
reduction takes place
(RED CAT)
Differences:
Voltaic Cell
Electrolytic Cell
Chemical reactions
produce electricity
Electricity produces
chemical reactions
Anode is negative
Cathode is positive
Anode is positive
Cathode is negative