1. Chapter 21
Electrochemistry

2. Electrochemical Processes
Chemical processes can either
release energy or absorb energy.
The energy can sometimes be in
the form of electricity.
Electrochemistry has many
applications in the home as well as in industry.
Flashlight and automobile batteries are examples
of devices used to generate electricity.
Biological systems also use electrochemistry to
carry nerve impulses.

3. Spontaneous Redox Reaction
When a strip of zinc metal is dipped into an
aqueous solution of blue copper sulfate, the zinc
becomes copper-plated.
Zn(s) + Cu2+(aq) Zn2+ (aq) + Cu(s)
The net ionic equation involves only zinc and
copper.
Electrons are transferred from zinc atoms to copper
ions.
This is a redox reaction that occurs spontaneously

4. Spontaneous Redox Reaction
As the reactions proceeds, zinc atoms lose
electrons as they are oxidized to zinc ions.
The zinc metal slowly dissolves.
0 +2 +2 0
Zn(s) + Cu2+(aq) Zn2+ (aq) + Cu(s)
At the same time, copper ions in solution gain the
electrons lost by the zinc.
They are reduced to copper atoms are deposited
as metallic copper.

5. Spontaneous Redox Reaction
As the copper ions in solution are gradually replace
by zinc ions, the blue color of the solution fades.
Oxidation: Zn(s) Zn2+ (aq) + 2e-
Reduction: Cu2+(aq) + 2e- Cu(s)

6. Activity Series of Metals
Zinc is higher on
the list than
copper.
For any two
metals in an
activity series,
the more active
metal is the
more readily
oxidized.

7. Electric Current
When zinc is dipped into a copper sulfate solution,
zinc becomes plated with copper.
In contrast, when a copper strip is dipped into a
solution of zinc sulfate, the copper does not
spontaneously become zinc-plated.
This is because copper metal is not oxidized by zinc
ions.
When a zinc strip is dipped into a copper sulfate
solution, electrons are transferred from zinc metal
to copper ions.
This flow of electrons is an electric current.

8. Electrochemical Process
The zinc-metal–copper-ion system is an example of
the conversion of chemical energy into electrical
energy.
Electrochemical process – any conversion between
chemical energy and electrical energy
All electrochemical processes involve redox
reactions.
If a redox reaction is to be used a a source of
electrical energy, the two half-reactions must be
physically separated.

9. Electrochemical Cell
The electrons released by zinc must pass through an
external circuit to reach the copper ions if useful
electrical energy is to be produced.
In this case the system serves as an electrochemical
cell.
Also, an electric current can be used to produce a
chemical change.
That system, too, serves as an electrochemical cell.
Electrochemical cell – any device that converts
chemical energy into electrical energy or vice versa.

10. Voltaic Cells
In 1800, Italian physicist Alessandro Volta build the
first electrochemical cell that could be used to
generate a dire electric current.
Voltaic cells – are electrochemical cells used to
convert chemical energy into electrical energy.
Electrical energy is
produced in a voltaic cell
by spontaneous redox
reactions within the cell.

11. Constructing a Voltaic Cell
Half cell – one part of a voltaic cell in which either
oxidation or reduction occurs.
Typical half-cell consists of a piece of metal
immersed in a solution of its ions.
Example: one half-cell is a zinc rod immersed in a
solution of zinc sulfate.
Other half-cell is a copper rod immersed in a solution
of copper sulfate.
Half cells are connected by a salt bridge which is a tube
containing a strong electrolyte, often potassium sulfate.

12. Constructing a Voltaic Cell
A porous plate my be used instead of a salt bridge
The porous plate allows ions to pass from on half-cell to
the other but prevents the solutions from mixing
completely.
A wire carries the electrons in the external circuit from the
zinc rod to the copper rod.
A voltmeter or light bulb can be connected in the circuit.
The driving force of such a voltaic cell is the spontaneous
redox reaction between zinc metal and copper(II) ions
in solution.

13. Constructing a Voltaic Cell
The zinc and copper rods in this voltaic cell are the
electrodes.
Electrode – a conductor in a circuit that carries
electrons to or from a substance other than a
metal.
The reaction at the electrode determines whether the
electrode is labeled as an anode or a cathode
Anode – the electrode at which oxidation occurs
Cathode – the electrode at which reduction occurs.

14. Constructing a Voltaic Cell
Electrons are consumed at the cathode and its
labeled the positive electrode. (reduction occurs)
Electrons are produced at the anode and its labeled
the negative electrode. (oxidation occurs)
The reaction at the electrode determines whether the
electrode is labeled as an anode or a cathode
All parts of the voltaic cell remain balanced in terms
of charge at all times.
The moving electrons balance any charge that might
build up as oxidation and reduction occur.

16. Questions
At which electrode does oxidation take place?
At the anode (negative electrode)
Where does reduction take place?
At the cathode (positive electrode)
What path do the electrons given up by zinc follow?
They go through the wire and the electric light to the
copper electrode.
What happen to the electrons at the copper electrode?
They reduce copper ions to copper.

17. How a Voltaic Cell Works
These steps actually occur at the same time.
Zn(s) Zn 2+(aq) + 2e-
1. Electrons are produced at the zinc rod according to
the oxidation half reaction
2. The electrons leave the zinc anode and pass
through the external circuit to the copper rod.
3. Electrons enter the copper rod and interact with
copper ions in solution.
Cu 2+ (aq) + 2e- Cu(s)

18. How a Voltaic Cell Works
4. To complete the circuit, both positive and negative
ions move through the aqueous solutions via the
salt bridge.
The overall cell reaction (note the electrons in the
overall reaction must cancel.
Zn(s) Zn2+(aq) + 2e-
Cu2+ (aq) + 2e- Cu(s)
Zn(s) + Cu2+ (aq) Zn2+(aq) + Cu(s)

19. The Need for a Salt Bridge
As zinc is oxidized at the anode, Zn2+ ions enter the
solution and they have no negative ions to balance
their charges.
So a positive charge tends to build up around the
anode.
At the cathode, Cu2+ ions are reduced to Cu and
taken out of the solution leaving behind unbalanced
negative ions.
Thus, a negative charge tends to develop around the
cathode.

20. The Need for a Salt Bridge
The salt bridge allows negative ions, such as SO 42-, to
be drawn to the anode to balance the growing
positive charge.
Positive ions, such as K+, are drawn from the salt
bridge to balance the growing negative charge at
the cathode.

21. Representing Electrochemical Cells
You can represent the zinc-copper voltaic cell by
suing the following shorthand form.
Zn(s) ZnSO4(aq) CuSO4(aq) Cu(s)
The single vertical lines indicate boundaries of
phases that are in contact.
The double vertical lines represent the salt bridge or
porous partition that separates the anode
compartment from the cathode compartment.
The half-cell that undergoes oxidation (the anode) is
written first.

22. Using Voltaic Cells
as Energy Sources
The zinc-copper voltaic cell is no longer used
commercially.
Current technologies that use electrochemical
processes to produce electrical energy include dry
cells, lead storage batteries and fuel cells.
Dry Cells – a voltaic cell in which the electrolyte is a
paste.
Dry cells used when a compact, portable electrical
energy source is required.

23. Dry Cells
A common type of dry cell is a flashlight battery.
A zinc container is filled with a thick, moist electrolyte
paste of manganese (IV) oxide, zinc chloride,
ammonium chloride, and water.
A graphite rode is embedded in the paste.
The zinc container is the anode and the graphite rod
is the cathode.
The thick paste and its surrounding paper liner
prevent the contents of the cell from freely mixing,
so a salt bridge is not needed.

24. Dry Cells
Oxidation:
Zn (s) Zn2+ (aq) + 2e-
Reduction:
2MnO2(s) + 2NH4+(aq) + 2e-
Mn2O3(s) + 2NH3(aq) + H2O(l)

25. Dry Cells
In an ordinary dry cell, the graphite rod serves only as
a conductor and does not undergo reduction.
MnO2 is the species that is actually reduced.
The electrical potential of this cell starts out at 1.5V
but decreases steadily during use to about 0.8V.
Dry cells of this type are not rechargeable because
the cathode reaction is not reversible.

26. Alkaline Battery
The alkaline battery is an improved dry cell used for
the same purposes.
In the alkaline battery, the reactions are similar to
those in the common dry cell, but the electrolyte s a
basic KOH past.
This change in design eliminates the buildup of
ammonia gas and maintains the Zn electrode,
which corrodes more slowly under alkaline
conditions.

27. Alkaline Battery

28. Lead Storage Batteries
People depend on lead storage batteries to start their
cars.
Battery is a group of cells
connected together.
A 12-V car battery consists of
six voltaic cells connected
together.
Each cell produces about 2 V
and consists of lead grids.

29. Lead Storage Batteries
One set of grids, the anode, is packed with spongy
lead.
The other set, the cathode, is packed with lead(IV)
oxide.
The electrolyte for both half-cells in a lead storage
batter is concentrated sulfuric acid.
Using the same electrolyte for both half-cells allows
the cell to operate without a salt bridge or porous
separator.

30. Lead Storage Batteries
Pb(s) + SO42-(aq) PbSO4(s) + 2e-
PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4(s) + 2H2O(l)
When a lead storage battery discharges, it produces the
electrical energy needed to start a car.
The overall spontaneous redox reaction that occurs is the
sum of the oxidation and reduction half-reactions.
Pb(s) + PbO2(s) + 2H2SO4 (aq) 2PbSO4(s) + 2H2O(l)
The equation shows that lead sulfate forms during
discharge.

31. Lead Storage Batteries
Pb(s) + PbO2(s) + 2H2SO4 (aq) 2PbSO4(s) + 2H2O(l)
The sulfate slowly builds up on the plates, and the
concentration of sulfuric acid decreases.
2PbSO4(s) + 2H2O(l) Pb(s) + PbO2(s) + 2H2SO4 (aq)
The reverse reaction occurs when a lead storage battery
is recharged. This occurs when the car’s generator is
working properly.
The reverse reaction is not a spontaneous reaction. A
direct current must pass through the cell in a direction
opposite that of the current flow during discharge.

32. Lead Storage Batteries
In theory, a lead storage battery can be discharged and
recharged indefinitely, but in practice its lifespan is
limited.
Small amounts of lead sulfate fall from the electrodes and
collect on the bottom of the cell.
Eventually, the electrodes lose so much lead sulfate that
the recharging process is ineffective or the cell is
shorted out.

33. Lead Storage Battery

34. Fuel Cells
To overcome the disadvantages associated with lead
storage batteries, cells with renewable electrodes
have been developed.
Fuel cells – are voltaic cells in which a fuel
substance undergoes oxidation and from which
electrical energy is continuously obtained.
Fuel cells do not have to be recharged.
They can be designed to emit no air pollutants and to
operate more quietly and more const-effectively
than a conventional electrical generator

35. Fuel Cells
Simplest fuel cell is the hydrogen-0xygen fuel cell.
Fuel cells – are voltaic cells in which a fuel
substance undergoes oxidation and from which
electrical energy is continuously obtained.
Fuel cells do not have to be recharged.
They can be designed to emit no air pollutants and to
operate more quietly and more const-effectively
than a conventional electrical generator
There are three compartment separated from one
another by two electrodes made of porous carbon.

36. Fuel Cells
Oxygen (the oxidizer) is fed into the cathode
compartment
Hydrogen (the fuel) is fed into the anode
compartment.
The gases diffuse slowly through the electrodes.
The electrolyte in the central compartment is a hot,
concentrated solution of potassium hydroxide.
Electrons from the oxidation half-reaction at the
anode pass through an external circuit to enter the
reduction half reaction at the cathode.

37. Fuel Cells

38. Fuel Cells
Oxidation: 2H2(g) + 4OH-(aq) 4H2O(l) + 4e-
Reduction: O2(g) + 2H2O(l) + 4e- 4 OH-(aq)
The overall reaction in the hydrogen-oxygen fuel cell
is the oxidation of hydrogen to form water.
2H2 (g) + O2 (g) 2H2O (l)
Fuel cells were developed for space travel where
lightweight, reliable power systems are needed.
Fuel cells different from lead storage batteries in that
they are not self-contained.

39. Fuel Cells
Operation depends on a steady flow of fuel and oxygen
into the cell (where combustion takes place) and the
flow of the combustion product out of the cell.
In the case of the hydrogen fuel cell, the product is pure
water.
Both the electricity generated and the water produced are
consumed in space flights.
Fuel cells convert 75% of the available energy into
electricity.
Conventional electric power plant converts from 35% to
40% of the energy of coal to electricity.

40. Fuel Cells
Other fuels, such as methane (CH4) and ammonia
(NH3), can be used in place of hydrogen.
4NH3(g) + 3O2 (g) 2N2 (g) + 6H2O (g)
CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
Other oxidizers, such as chlorine (Cl2) and ozone
(O3)can be used in place of oxygen.

41. Question
Make a sketch of a tin/lead voltaic cell
Sn SnSO4 PbSO4 Pb
Label the cathode and anode, and indicate the
direction of electron flow. Write the equations for
the half-reactions.
Sn(s) Sn2+(aq) + 2e-
Pb2+(aq) + 2e- Pb(s)
Tin is the anode, lead is the cathode.
The electrons flow from tin to lead.

42. Questions
What type of reaction occurs during an
electrochemical process?
Redox
What is the source of electrical energy produced in a
voltaic cell?
Spontaneous redox reactions within the cell
If the relative activities of two metals are known,
which metal is more easily oxidized?
The metal with the higher activity.

43. Homework
Draw detailed pictures of the following:
1. A Voltaic Cell
2. A Lead Storage Battery
3. Fuel Cell
• Make sure to include all substances involved as
well as the half reactions and overall reaction.
• Make sure you label the anode and cathode and
indicate where oxidation and reduction take place.
• List common uses of each
• List advantages of disadvantages of each.

44. End of Section 20.1

45. Electrical Potential
Electrical potential of a voltaic cell is a measure of the
cell’s ability to produce an electric current. (in Volts)
You cannot measure the electrical potential of an
isolated half-cell.
When two half-cells are connected to form a voltaic
cell, the difference in potential can be measured.
The electrical potential of a cell results from a
competition for electrons between two half-cells.
The half-cell that has a greater tendency to acquire
electrons will be the one in which reduction occurs.

46. Electrical Potential
Oxidation occurs in the other half-cell as there is a
loss of electrons.
Reduction potential – the tendency of a given half-
reaction to occur as a reduction.
The half-cell in which reduction occurs has a greater
reduction potential than the half-cell in which
oxidation occurs.
Cell potential – the difference between the reduction
potentials of the two half-cells.

47. Cell Potential
reduction potential reduction potential
Cell Potential = of half-cell in which - of half-cell in which
reduction occurs oxidation occurs
E0cell = E0red - E0oxid
Standard cell potential (E0cell) – is the measured cell
potential when the ion concentration in the half-cells are
1M, any gases are at a pressures of 101 kPa, ant the
temperature is 25ºC.
The symbols E0red and E0oxid represent the standard
reduction potential for the reduction and oxidation half-
cells, respectively.

48. The Lemon Battery
A working voltaic cell made
using a lemon and strips of
copper and zinc.
Which is the anode and which
is the cathode?
Zn is anode & Cu is cathode
What process goes on at the anode and cathode?
Oxidation at anode, reduction at cathode.
Which process are electrons lost?
Oxidation
What role does the lemon play in the battery?
Salt bridge

49. Standard Hydrogen Electrode
Because half-cell potentials cannot be measured
directly, scientists have chosen an arbitrary
electrode to serve as a reference.
The standard hydrogen electrode is used with other
electrodes so the reduction potentials of the other
cells can be measured.
The standard reduction potential of the hydrogen
electrode has been assigned a value of 0.00 V.

50. Standard Hydrogen Electrode
Consists of a platinum electrode
immersed in a solution with a
hydrogen-ion concentration of 1M
Solution is at 25 C and the electrode
is a small square of platinum foil
coated with finely divided platinum,
known as platinum black.
Hydrogen gas at 101 kPa is bubbled
around the platinum electrode.

51. Standard Hydrogen Electrode
2H+ (aq, 1M) + 2e- H2 (g, 101kPa) EOH+ = 0.00V
Double arrows indicate the reaction is reversible.
Standard reduction potential of H+ is the tendency
of H+ ions to acquire electrons and be reduced to
H2 (g)
Whether the half-cell reaction occurs as a reduction
or as an oxidation is determined by the reduction
potential of the half-cell to which the standard
hydrogen electrode is connected.

52. Standard Reduction Potentials
A voltaic cell can be made by connecting a standard
hydrogen half-cell to a standard zinc half-cell.

53. Standard Reduction Potentials
To determine the overall reaction for this cell, first
identify the half-cell in which reduction takes place.
In all electrochemical cells, reduction takes place at
the cathode and oxidation takes place at the anode.
A voltmeter gibes a reading of +0.76 V when the zinc
electrode is connect to the negative terminal and
the hydrogen electrode is connect to the positive
terminal.
Zinc is oxidized – anode and Hydrogen ions are
reduced – hydrogen electrode is the cathode. .

54. Standard Reduction Potentials
Oxidation: Zn (s) Zn2 (aq) + 2e- (at anode)
Reduction: 2H+(aq) + 2e- H2 (g) (at cathode)
You can determine the standard reduction potential of
a half-cell by using a standard hydrogen electrode
and the equation for standard cell potential
E0cell = E0red - E0oxid
E0cell = E0H+ - E0Zn2+
0.76 V = 0.00 V - E0Zn2+
E0Zn2+ = -0.76 V

55. Standard Reduction Potentials
The standard reduction potential for the zinc half-cell
is -0.76 V.
The value is negative because the tendency of zinc
ions to be reduced to zinc metal in this cell is less
than the tendency of hydrogen ions to be reduced
to hydrogen gas.
Consequently, the zinc ions are not reduced. Instead,
the opposite occurs: Zinc metal is oxidized to zinc
ions.

56. Standard Reduction Potentials
For a standard copper half-cell, the measured
standard cell potential is +0.34V.
Copper is the cathode and Cu2+ ions are reduced to
Cu metal
Hydrogen half-cell is the anode, and H2 gas is
oxidized to H+ ions.
E0cell = E0red - E0oxid
E0cell = E0Cu2+ - E0H+
+0.34 V = E0Cu2+ - 0.00
E0Cu2+ = +0.34 V

57. Discussion
The two half-cells of a voltaic cell are competing for
electrons.
Oxidation or reduction could occur in either cell.
The half-cell with the more positive reduction potential
will win the competition and undergo reduction.
The potential produced by the electrochemical cell is
the difference in the reduction potentials of the two
half-cell reactions.
The quantitative value of any half-cell potential is
obtained by measuring it against the standard
hydrogen electrode. (Review table 21.2 page 674)

58. Discussion
Which reactions have the greatest tendency to occur
as reductions?
Activity series of metals have the most active metals
at the top. Because active metals lose electrons
easily, they are most easily oxidized.
Thus, the ions of active metals are least likely to be
reduced.
The potential produced by the electrochemical cell is
the difference in the reduction potentials of the two
half-cell reactions.

59. Calculating Standard Cell Potentials
To function, a cell must be constructed of two half-
cells.
The half-cell reaction having the more positive (or
less negative) reduction potential occurs as a
reduction in the cell.
You can use the know standard reduction potentials
for various half-cells to predict the half-cell in which
reduction and oxidation will occur.
If the cell potential for a given redox reaction is + then
the reaction is spontaneous as written. If the cell
potential is - , then the reaction in nonspontaneous.

60. Question
Determine whether the following redox reaction will
occur spontaneously.
3Zn2+(aq) + 2Cr(s) 3Zn(s) + 2Cr3+(aq)
Oxidation: Cr(s) Cr3+(aq) + 3e- E0Cr3+ = -0.74V
Reduction: Zn2+(aq) + 2e- Zn(s) E0Zn2+ = -0.76V
E0cell = E0red - E0oxid
E0cell = E0Zn2+ - E0Cr3+
E0cell = -0.76 – (-0.74)
E0cell = -0.02 V (nonspontaneous)

61. Question
Is this redox reaction spontaneous as written?
Co2+(aq) + Fe(s) Co(s) + Fe2+(aq)
Oxidation: Fe(s) Fe2+(aq) + 2e- E0Cr3+ = -0.44V
Reduction: Co2+(aq) + 2e- Co(s) E0Zn2+ = -0.28V
E0cell = E0red - E0oxid
E0cell = E0Co2+ - E0Fe2+
E0cell = -0.28 – (-0.44)
E0cell = +0.16 V (spontaneous)

62. Question
Determine the cell reaction for a voltaic cell
composed of the following half-cells.
Fe3+(aq) + e- Fe2+(aq) E0Fe3+ = +0.77V
Ni2+(aq) + 2e- Ni(s) E0Ni2+ = -0.25V
The half-cell with the more positive reduction potential
is the one in which reduction occurs (the cathode)
Oxidation: Ni(s) Ni2+(aq) + 2e-
Reduction: 2Fe3+(aq) + 2e- 2Fe2+(aq) (balance e-)
Ni(s) + 2Fe3+(aq) Ni2+(aq) + 2Fe2+(aq)

63. Discussion
In the previous example, we had to multiply the Fe
half-cell reaction by a factor of 2 to cancel out the
electrons.
Even though, there were two times as many electrons
present, the tendency for the electrons to flow is not
two times greater.
The tendency, which is measured by the E0 value,
remains the same.

64. Questions
A voltaic cell is constructed using the following half
reactions.
Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V
Al3+(aq) + 3e- Al(s) E0Al3+ = -1.66V
2Al(s) + 3Cu2+(aq) 2Al3+(aq) + 3Cu(s)
A voltaic cell is constructed using the following half
reactions.
Ag+(aq) + e- Ag(s) E0Ag+ = +0.80V
Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V
Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

65. Question
Calculate the standard cell potential for the Ni/Fe
voltaic cell. Half-reactions are as follows;
Fe3+(aq) + e- Fe2+(aq) E0Fe3+ = +0.77V
Ni2+(aq) + 2e- Ni(s) E0Ni2+ = -0.25V
E0cell = E0Fe3+ - E0Ni2+
E0cell = +0.77 V – (-0.25 V)
E0cell = +1.02 V

66. Question
A voltaic cell is constructed using the following half-
reactions
Al3+(aq) + 3e- Al(s) E0Al3+ = -1.66V
Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V
E0cell = E0Cu2+ - E0Al3+
E0cell = +0.34 V – (-1.66 V)
E0cell = +2.00 V

67. Question
A voltaic cell is constructed using the following half-
reactions
Ag+(aq) + e- Ag(s) E0Ag+ = +0.80V
Cu2+(aq) + 2e- Cu(s) E0Cu2+ = +0.34V
E0cell = E0Ag+ - E0Cu2+
E0cell = +0.80 V – (+0.34 V)
E0cell = +0.46 V

68. Questions
What causes the electrical potential of a cell?
Competition for electrons between two half-cells
What is the electrical potential of a standard hydrogen
electrode?
Assigned a value of 0.00 V at 25ºC
How can you find the standard reduction potential of a
half-cell?
By connecting it to a standard hydrogen electrode and
measuring the cell potential
What cell potential values indicate a spontaneous
reaction? A nonspontaneous reaction?
Positive cell potential - spontaneous

69. Homework
Using the reduction potentials from table 21.2, create an
electrochemical cell that will operate spontaneously.
Calculate the cell potential
Write the equations for the two half-reactions and the
overall cell reaction.
Use the shorthand method to represent the cell

70. End of section 20.2

71. Electrolytic Cells
An electric current can be used to make a
nonspontaneous redox reaction go forward.
Electrolysis – the process in which electrical
energy is used to bring about a nonspontaneous
chemical change.
Examples of electrolysis are silver-plated dishes and
utensils, gold-plated jewelry, and chrome-plated
automobile parts.
Electrolytic cell – the apparatus in which
electrolysis is carried out is an electrochemical
cell used to cause a chemical change through the
application of electrical energy.

72. Electrolytic Cells
An electrolytic cell uses electrical energy (direct current)
to make a nonspontaneous redox reaction proceed to
completion.
In both voltaic and electrolytic cells, electrons flow from
the anode to the cathode in the external circuit.
For both types of cells, the electrode at which reduction
occurs is the cathode.
The key difference between voltaic and electrolytic cells is
that in a voltaic cell, the flow of electrons is the result of
a spontaneous redox reaction, whereas in an
electrolytic cell, electrons are pushed by an outside
power source, such as a battery.

73. Voltaic cell – energy is released from a spontaneous
redox reaction.
Electrolytic cell - energy is absorbed to drive a non-
spontaneous reaction.

74. Electrolytic Cells
Electrolytic and voltaic cells also differ in the
assignment of charge to the electrodes.
In an electrolytic cell, the cathode is considered to be
the negative electrode, because it is connected to
the negative electrode of the battery.
The anode in an electrolytic cell is considered to be
the positive electrode because it is connected to
the positive electrode of the battery.
In a voltaic cell, the anode is the negative electrode
and the cathode is the positive electrode.

75. Electrolytic Cells
Electrolytic processes are used to separate active
metals such as aluminum, magnesium, and sodium
from their salts.
The same process is used to recover metals from
ores.

76. Electrolysis of Water
When a current is applied to two electrodes immersed
in pure water, nothing happens.
When an electrolyte such as sulfuric acid or
potassium nitrate in low concentration is added to
the pure water, the solution conducts electricity and
electrolysis occurs.
The products of the electrolysis of water are hydrogen
gas and oxygen gas.

77. Electrolysis of Water
Water is reduced to hydrogen at the cathode
Reduction: 2H2O(l) + 2e- H2(g) + 2OH-(aq)
Water is oxidized at the anode
Oxidation: 2H2O(l) O2(g) + 4H+ (aq) + 4e-
The region around the anode turns acidic due to an
increase in H+ ions.
The region around the cathode turns basic due to the
production of OH- ions.

78. Electrolysis of Water
The overall cell reaction
4H2O(l) + 4e- 2H2(g) + 4OH-(aq) (x2 to balance)
2H2O(l) O2(g) + 4H+ (aq) + 4e-
6H2O(l) 2H2(g) + 4OH-(aq) + O2(g) + 4H+(aq)
The ions produced tend to recombine to form water,
so they are not included in the net reaction.
6H2O(l) electrolysis
H2(g) + O2(g)

79. Electrolysis of Brine
If the electrolyte in an aqueous solution is more easily
oxidized or reduced than water, then the products of
electrolysis will be substances other than hydrogen and
oxygen.
Example is brine (a concentrated aqueous solution of
sodium chloride) which produces chlorine gas,
hydrogen gas, and sodium hydroxide.
During electrolysis of brine, chloride ions are oxidized to
produce chlorine gas at the anode.
Oxidation: 2Cl-(aq) Cl2(g) + 2e- (at anode)

80. Electrolysis of Brine
Water is reduced to produce hydrogen gas at the
cathode.
Reduction: 2H2O(l) + 2e- H2(g) + 2OH-(aq) (at cathode)
Sodium ions are not reduced to sodium metal in the
process because water molecules are more easily
reduced than are sodium ions.
The reduction of water produces hydroxide ions as well
as hydrogen gas. Thus the electrolyte in solution
becomes sodium hydroxide.

81. Electrolysis of Brine
The overall ionic equation
2H2O(l) + 2e- H2(g) + 2OH-(aq)
2Cl-(aq) Cl2(g) + 2e-
2H2O(l) + 2Cl-(aq) H2(g) + 2OH-(aq) Cl2(g)
The spectator ion Na+ can be included in the equation (as
part of NaCl and of NaOH) to show the formation of
sodium hydroxide during the electrolytic process
2NaCl (aq) + 2H2O(l) + 2Cl-(aq) H2(g) + 2NaOH(aq) Cl2(g)

82. Electrolysis in Metal Processing
Electrolytic cells are commonly used in the plating,
purifying and refining of metals.
Many of the shiny, metallic objects you see every day,
such as chrome-plated fixtures or nickel-plated coins,
were manufactured with the help of electrolytic
processes.
Electroplating is the deposition of a think layer of metal
on an object in an electrolytic cell.
An object may be electroplated to protect the surface of
the base metal from corrosion or to make it more
attractive.

83. Electrolysis in Metal Processing
An object that is to be silver-plated is made the cathode in
an electrolytic cell.
The anode is the metallic silver that is to be deposited
The electrolyte is a solution of a silver salt, such as silver
cyanide.
When a direct current is applied, silver ions move from
the anode to the object to be plated.
Reduction: Ag+ (aq) + e- Ag (s) (at cathode)
The net result is that silver transfers from the silver
electrode to the object being plated.

84. Electrolysis in Metal Processing
Many factors contribute to the quality of the metal
coating that forms.
In the plating solution, the concentration of the
cations to be reduced must be carefully controlled.
The solution must also contain compounds to control
the acidity and to increase the conductivity.
Other compounds may be used to make the metal
coating brighter or smoother.

85. Electrolysis in Metal Processing
Electroforming – is a process in which an object is
reproduced by making a metal mold of it at the
cathode of a cell.
A phonograph record can be coated with metal so it
will conduct a current.
It is then electroplated with a thick coating of metal.
This coating can be stripped off and used as a mold
to produce copies of the record.

86. Electrowinning
Electrowinning – a process where impure metals
can be purified in electrolytic cells.
The cations of molten salts or aqueous solutions are
reduced at the cathode to give very pure metals.
A common use is in the extraction of aluminum form
its ore, bauxite. (Al2O3)
In a method know as the Hall-Heroult process, purified
alumina is dissolved in molten cryolite (Na3AlF6), and
heated to above 1000ºC in a carbon line tank.

87. Electrowinning
The carbon lining, connected to a direct current,
serves as the cathode. The anode consists of
carbon rods dipped into the tank.
At the cathode, Al3+ ions are reduced, forming molten
aluminum. At the anode, carbon is oxidized,
forming carbon dioxide gas.
2Al2O3(l) + 3C(s) 4Al(l) + 3CO2(g)

88. Other Electrolytic Processes
Electrorefining – a piece of impure metal is made the
anode of the cell. It is oxidized to the cation and then
reduced to the pure metal at the cathode.
Electrorefining technique is used to obtain ultrapure silver,
lead and copper.
Other electrolytic processes are centered on the anode
rather than the cathode.
Electropolishing - the surface of an object at the anode
is dissolved selectively to give it a high polish.
Electromachining - a piece of metal at the anode is
partially dissolved until the remaining portion is an exact
copy of the object at the cathode.

89. Questions
What is the difference between an electrolytic cell and a
voltaic cell?
Voltaic cell uses an electrochemical reaction to produce
electrical energy. An electrolytic cell uses electrical
energy to bring about a chemical change.
What products form during the electrolysis of water?
H2 (g) and O2 (g)
What chemical changes occur during the electrolysis of
brine?
Chloride ions are oxidized to produce chlorine gas and
water is reduced to produce hydrogen gas.

90. Questions
What are some application of electrolysis in the field of
metallurgy?
Electroplating (deposition of a thin layer of metal on an
object), electrorefining (purification of metals ) and
electrowinning (extraction of metals)
What is the charge on the anode of an electrolytic cell? Of
a voltaic cell?
Electrolytic cell anode (+); voltaic cell anode (-)
Which process, oxidation or reduction, always occurs at
the cathode of an electrolytic cell?
Reduction

91. Questions
Can metallic sodium be obtained by electrolyzing brine?
No; the products are chlorine gas, hydrogen gas, and
sodium hydroxide.
Sodium is obtained by electrolysis of molten NaCl in the
Downs cell, which operates at 801ºC to keep it melted.
The anode and cathode of the cell are separated to
prevent recombination of sodium and chlorine.
Reduction of Na+ occurs at a graphite anode. Liquid
sodium rises to the top of the molten NaCl and is drawn
off.
2NaCl (l) 2Na (l) + Cl2 (g)

92. End of Chapter 20