Batteries and Electrolytic Cells PowerPoint (pdf)

BATTERIES AND ELECTROLYTIC
CELLS
Practical Electrochemistry

How Batteries Work
• One of the most practical
applications of spontaneous redox
reactions is making batteries.
• In a battery, a spontaneous
electron transfer from a reducer to
an oxidizer is forced to occur
through an external wire.
• As the electrons crowd through
the wire, they give up their energy
to do work.

Spontaneous Electron Transfer
• Zinc atoms will give up their 2
valence electrons to copper ions
easily.
2+
Zn
Cu 0 (s)
Cu 2+

Remote Transfer
• If a barrier is placed between Zn and
Cu 2+ , the transfer can still take place,
but through a wire.
Zn

Electrochemical Cells
• Electrochemical cells are those in
which spontaneous electron
transfer takes place through a
wire.
• They are also called voltaic or
galvanic cells.
• You probably know them best by
their practical name: batteries.
• We think of batteries as electrical,
but they are really chemical
devices that produce electricity.

Batteries
• Batteries are really quite simple
devices that contain 2 half-cells
with electrodes and solutions.
• Each half-cell is usually made of a
metal electrode and a solution
containing a salt of the same
metal.
• There is an external wire
connecting the 2 electrodes.
• The electrodes are the negative
anode and the positive cathode.

Batteries
• An electrical power using device
(resistor) is placed in the wire, and
as the electrons crowd through the
device, they give up their energy
to make it operate: ex, a light bulb
in a flashlight.
• The circuit must however be
complete for electrons to flow, so
a salt bridge or porous barrier
between half-cells is also
necessary.

Basic Battery
Zinc metal
atoms with 2
electrons
Copper
atoms with 2
electrons
Zn 2+
aqueous
ions
2+
Cl-
2+
Hungry Cu
2+ ions
Cl-
Zn
Cu
Anions move
across the
2+ Cl- Cl-
barrier to
Cl-
Cl-
complete the
circuit.
Anode half-cell (-) Cathode half-cell (+)
Zn Zn 2+ + 2e - Cu 2+ + 2 e - Cu

Disclaimer
• Unlike the animations, electrons
do not really “zip” through a wire
from one end to the other.
• It’s actually more like a tube full of
ping-pong balls.
• If you push a new ball into one
end, they all move down a little,
and the one on the other end pops
out.
• Electrons in a real electric circuit
actually move quite slowly.

Basic Battery
• The anode is slowly dissolving
away as the atoms become
soluble ions.
• The cations of the cathode
compartment are becoming metal
atoms, so the cathode is getting
more massive.
• When either the anode is
completely gone or the cathode
cations are gone, the battery is
dead.

Why batteries die on the shelf
2+
2+
Cl-
Cl-
Cl-
Cl-
- Cl- +
Zn
Cu
Cl-
All the energy is released as heat and no work is
done. However, the barrier must be porous for the
battery to function normally.

Rechargeable Batteries
Zn
2+
2+
Generator:
electron pump
Cl-
2+
Cl-
Cl-
Cl-
Cu
Cl-
Cl-
Cl-
Cl-
With a
rechargeable
battery, a
generator pushes
the electrons
backwards in the
circuit.
Even though the
electrons don’t
“want” to go in
this direction,
the generator
forces them.
The electrons
are restored to
the anode and
the battery can
be used again.

Electromotive Force (emf)
• What is electromotive force?
• Emf is symbolized E°.
• It is the “pressure” with which
electrons flow through a circuit.
• Emf is measured in volts.
• So, how does one go about
measuring the voltage produced
by a battery?
• Or, are all batteries the same
except for size?

Half-Cell Potentials
• Years ago, scientists agreed on a
standard reference half-cell, against
which to measure all other halfreactions.
• The reference they chose was a halfcell
with H + in the solution, and a
platinum electrode.
• The reaction occurring in the cell is:
2H + (aq) + 2 e - H 2 (g)
• Hydrogen ions are reduced.
• (The Pt is an inert electrode.)

Half-Cell Potentials
• The hydrogen half-cell was assigned a
voltage of E° = 0.0Volts.
• Various reactions were used for the
other half-cell and the voltages
recorded.
• If an element is easier to reduce than
hydrogen, the voltage is a positive
value.
• For example, when the hydrogen half
cell is hooked up to a copper half-cell,
the voltage is +0.34V.

Half-Cell Potentials
• So, a copper/hydrogen battery has
an overall voltage (E° cell ) of +0.34
volts.
• If the element is more difficult to
reduce than hydrogen, the
recorded voltage is negative.
• Ex. Cr 3+ + 3e - Cr E° = -0.74V
• All such half-cell voltages are
recorded in a table of Standard
Electrode Reduction Potentials.

Standard Reduction Potentials
Reduction Half-Reaction E°
F 2 + 2e - 2 F -
+2.87V
Au 3+ + 3e - Au
+1.42V
Ag + + e - Ag
+0.80V
I 2 + 2e - 2I -
+0.54V
Cu 2+ + 2e - Cu
+0.34V
2 H + + 2 e - H 2 (g) 0.0 0V
Pb 2+ + 2e - Pb
-0.13V
Ni 2+ + 2e - Ni
-0.28V
Cd 2+ + 2e - Cd
-0.41V
Zn 2+ + 2e - Zn
-0.76V
Li + + e - Li
-3.04V

Battery Voltages
• In order for an electrochemical cell
to work as a battery, the overall
cell voltage must be positive.
• Remember that a reduction cannot
occur without an oxidation.
• The table we just saw shows only
reductions.
• So, one of the half-cell reactions
will need to be turned around to
become an oxidation.

Battery Voltages
• Guess what happens to the
symbol on the voltage if we turn
the equation around.
• The +/- symbol changes!
• Let’s examine our Cu/Zn battery
again with a look at the voltage it
generates.

Cu/Zn Battery
• Here are the 2 half-reactions from
the table:
• Cu 2+ + 2e - Cu +0.34V
• Zn 2+ + 2e - Zn -0.76V
• One of the reactions will have to be
turned around to show an oxidation.
• Which one will it be to leave a positive
voltage?
• Right! The Zn reaction must be
reversed.

Cu/Zn Battery
• Cu 2+ + 2e - Cu
• Zn Zn 2+ + 2e -
+0.34V
+0.76V
• Notice that the voltage became
positive.
• To calculate battery voltage, simply
add the 2 equations together.
• Cu 2+ + 2e - Cu +0.34V
• Zn Zn 2+ + 2e - +0.76V
• Cu 2+ + Zn Zn 2+ + Cu +1.10V

Battery Electrodes
• In an earlier slide, we said that a
battery has an anode and a
cathode.
• How can you tell which is which?
• Do you know the legend of Paul
Bunyan?
• Paul, the giant lumberjack had a
well known pet: Babe, the blue ox.
• But Paul had another, less famous
pet: Fritz, the red cat.

Battery Electrodes
• So, when you think of
batteries, think of Paul
Bunyan.
• He had 2 pets: an ox and
a red cat.
• The anode is where
oxidation happens, and
reduction is at the
cathode.
• The cathode is the +
electrode, and the anode
is the – electrode.
• Electron flow is always
from anode to cathode.

Try This Example
• A battery is composed of a silver
half-cell and a cadmium half-cell.
• Write the reduction and oxidation
half reactions.
• Draw the battery.
• Calculate the voltage.
• Label the anode and cathode.
• Draw the direction of electron flow.

Silver-Cadmium Battery
e - flow:
Cd to Ag
Cd is
oxidized:
anode
Ag is
reduced:
cathode
Put in the porous
barrier.
Add the
electrodes.
Ag Cd Put cations in
Anode the solution.
-
Ag + Cd 2+ We also need
some anions.
Cathode
+
Write half-reactions and voltages.
Ag + + e - Ag E° = 0.80V
Cd Cd 2+ Cd + 2e 2+ - + 2eCd
-
E° = +0.41V -0.41V
Total Voltage = 1.21V
Connect the
external circuit.
Reverse the Cd
equation to be
an oxidation.

Alternate Battery Diagrams
• Sometimes a simple battery
diagram is shown as 2 separate
half-cells. Instead of a porous
barrier, there is a “salt bridge.”
The bridge is
usually some sort
of tube that is
filled with salt
solution. It allows
transfer of ions
from one half-cell
to the other.

Electrolysis
• Electrolysis involves using electric
current and energy to force a nonspontaneous
redox reaction to go.
• It is the opposite of an
electrochemical cell.
• So, the overall cell voltage (E°) is
negative rather than positive.
• This is the minimum voltage
needed to make the cell operate.

Electrolysis
• The rechargeable battery we saw
earlier is an example of an
electrolytic reaction.
• We use a direct current generator
for force the electron flow to go in
reverse.
• Electrolysis is also used
extensively for electroplating one
metal onto another.

Electrolytic Cell Example
• An easy example of an electrolysis
reaction would be the
decomposition of NaCl.
• First the NaCl must be melted
(around 1500°C).
• Then inert metal electrodes are
placed in the liquid and the current
is turned on.
• An electron is removed from Cl -
and transferred to Na + .

Electrolytic Cell Example
+ Na + ion
Na atom
At the anode, Cl - ions lose
electrons and are oxidized
to Cl 2
molecules.
- Cl - ion electrons
Cl atom
Electrons still flow
from anode to
cathode.
At the cathode, Na +
ions are reduced to
Na atoms.
+
- +
-
A
-
+ C

Electrolytic Cell
• The key differences
between the
electrolytic cell and
the electrochemical
cell are:
1. Generator rather
than a resistor in the
circuit.
2. Lack of a porous
barrier or salt bridge.
3. Polarities of the
electrodes are
reversed. The
anode is positive
and the cathode is
negative.
+
- +
-
A
-
+ C

Aluminum Production
• Another very important industrial
electrolysis reaction is the Hall-
Heroult process for making
aluminum from bauxite ore.
• Al 2 O 3 (l) + 3 C(s) 2 Al(l) + 3 CO(g)
• Al 3+ + 3 e - Al metal (reduction)
• C(s) + O CO (oxidation of C)
• The production of aluminum metal
takes about 1% of all the electricity
used in the U. S. each year.

Aluminum Production
• Although aluminum is relatively
inexpensive, (about $1.50-$2.00
per pound), it is MUCH cheaper to
melt and recycle aluminum than it
is to smelt new metal.
• In the early 1800’s, when only a
small amount of metal could be
produced, Al cost $250,000 per
pound ($545/gram) while silver
was only $17/gram!

Calculating Voltage
• Voltage for electrolytic cells is
calculated the same way that it is
for batteries, except that the value
will be negative rather than
positive for this non-spontaneous
cell.
• 2Cl - Cl 2 + 2 e - E° = -1.36V
• 2Na + + 2 e - 2Na E° = -2.71V
• Total voltage: E° cell = -4.07V
• 4.07V is the minimum voltage
needed to decompose NaCl(l).

Uses for Electrolysis
• Electrolysis of seawater is used to
make Cl 2 , H 2 and NaOH, three
very useful and important
materials.
• Electrolysis is also used
extensively for electroplating one
metal onto another (sometimes
called ‘anodizing.’)
• For example ‘silver’ flatware is
really silver plated flatware.

Electroplating
• Silverware is not solid silver.
• It is made of some cheaper,
harder metal, like nickel.
• Then the nickel piece is coated
with silver in an electrolysis cell.
• A bar of solid silver acts as the
anode, and the nickel piece is the
cathode.
• The cell simply transfers silver
atoms from anode to cathode.

Electroplating
Silver bar anode:
Ag atoms change
to ions and
dissolve into the
solution.
e -
The generator pulls
electrons from the
anode and pushes
them into the
cathode.
e -
e -
e -
e -
e - e - e -
e -
Knife and
threek
cathode: Ag +
ions change
back to atoms
and stick to, or
plate the
metal.
Ag +
Ag
anode
Ag
AgNO 3
(aq)

Schematic Symbols
• Here are some symbols commonly
used in electronic schematics:
• Generator
• Resistor
• Battery
• Voltmeter
V
• Ammeter A
• Switch