1. Dr. Floyd Beckford Lyon College
CHM 120 CHAPTER 21 Electrochemistry: Chemical Change and Electrical Work
Dr. Floyd Beckford
Lyon College

2. REVIEW Oxidation: the loss of electrons by a species,
leading to an increase in oxidation number of one
or more atoms
Reduction: the gain of electrons by a species,
leading to an decrease in oxidation number of one
Oxidizing agents: the species that is reduced in a
redox reaction

3. Reducing agents: the species that is oxidized in a redox reaction

4. In acidic solution: add H+ or H2O only
In basic solution: add OH- or H2O only
To balance O
To balance H
In acidic solution:
Add H2O
and then
Add H+
For each O needed
For each H needed
In basic solution
1. add 2 OH- to the side needing O and
1. add 1 H2O to the side needing H and
2. add 1 H2O to the other side
2. add 1 OH- to the other side

5. THE HALF-REACTION METHOD
This method breaks the overall reaction into its
two components – half-reactions
Each half-reaction is balanced separately and
then added
Use the following guidelines to help
Write as much of the unbalanced net ionic
equation as possible
2. Decide which atoms are oxidized and which are
reduce – write the two unbalanced half-reactions

6. 3. Balance by inspection all atoms in each half-
reaction except H and O
4. Use the rules mentioned previously to balance
H and O in each half-reaction
5. Make equal the number of electrons involved in
both half-reactions
Take a look at the breathalyzer reaction
H+(aq) + Cr2O72-(aq) + C2H5OH(l) 
Cr3+(aq) + C2H4O(l) + H2O(l)

7. Balance the following net ionic equation in basic
solution.

9. ELECTROCHEMISTRY Deals with chemical changes produced by an
electric current and with the production of
electricity by chemical reactions
All electrochemical reactions involve transfer
of electrons and are redox reactions
EChem reactions take place in electrochemical
cell (an apparatus that allows a reaction to
occur through an external conductor)

10. ELECTROCHEMICAL CELLS
Two types:
1. Electrolytic cells: - these are cells in which an
external electrical source forces a
nonspontaneous reaction to occur
2. Voltaic cells: - also called galvanic cells. In
these cells spontaneous chemical reactions
generate electrical energy and supply it to an
external circuit

11. Electric current enters and exits the cell by
electrodes - electrodes are surfaces upon which
oxidation or reduction half-reactions occur
Inert electrodes: - electrodes that don’t react
Two kinds of electrodes:
1. Cathode: - electrode at which reduction
occurs (electrons are gained by a species)
2. Anode: - electrode at which oxidation occurs
(as electrons are lost by some species)

12. VOLTAIC CELLS Cells in which spontaneous reactions produces
electrical energy
The two half-cells are separated so that electron
transfer occurs through an external circuit
Each half-cell contains the oxidized and reduced
forms of a species in contact with each other
Half-cells linked by a piece of wire and a salt
bridge

14. A salt bridge has three functions:
1. It allows electrical contact between the two
half-cells
2. It prevents mixing of the electrode solutions
3. It maintains electrical neutrality in each
half-cell as ions flow into and out of the salt
bridge
Point 2 is important – no current would flow if
if both solutions were in the same cell

15. Point 3 is also important – anions flow into the
oxidation half-cell to counter the build-up
of positive charge
Current flow spontaneously from negative to
the positive electrode
In all voltaic cells the anode is negative and the
cathode is positive

17. In voltaic cells, voltage drops as the reaction
proceeds. When voltage = 0, the reaction is at
equilibrium

18. The Silver-Copper cell
Composed of two half-cells:
1. A strip of copper immersed in 1 M CuSO4
2. A strip of silver immersed in 1 M AgNO3
Experimentally we see:
: - Initial voltage is 0.46 volts
: - The mass of the copper electrode decreases
: - The mass of the silver electrode increases
: - [Cu2+] increases and [Ag+] decreases

19. Cu  Cu2+ + 2e- (oxidation, anode)
2(Ag+ + e-  Ag) (reduction, cathode)
2Ag+ + Cu  Cu2+ + Ag (Overall cell reaction)
Cu |Cu2+(1.0 M) ||Ag+(1.0 M) | Ag
Notice that in this case the copper electrode is
the anode

20. STANDARD ELECTRODE POTENTIALS
Associated with each voltaic cell is a potential
difference called the cell potential, Ecell
E measures the spontaneity of the cell’s redox
reaction
Higher (more positive) cell potentials indicate a
greater driving force for the reaction as written
All electrode potentials are measured versus the
Standard Hydrogen Electrode (SHE): E° = 0.00 V

21. The E°cell calculated is for the cell operating
under standard state conditions
For electrochemical cell standard conditions
are:
solutes at 1 M concentrations
gases at 1 atm partial pressure
solids and liquids in pure form
All at some specified temperature, usually 298 K

23. The electrode potential for each half-reaction is
written as a reduction process
The more positive the E° value for a half-
reaction the greater the tendency for the reaction
to proceed as written
The more negative the E° value, the more likely
is the reverse of the reaction as written

24. Prediction of Spontaneity
1. First write the HR equation with the more
positive (less negative) E° for the reduction along
with its potential
2. Write the other HR as an oxidation and include
its oxidation potential
3. Balance the electron transfer
4. Add the reduction and oxidation HR and add
the corresponding electrode potentials to get the
overall cell potential, E°cell

25. Important points to note:
1. E° for oxidation half-reactions are equal to
but opposite in sign to reduction half-reactions
2. Half-reaction potentials are the same
regardless of the species’ stoichiometric
coefficient in the balanced equation
E°cell > Forward reaction is spontaneous
E°cell < Backward reaction is spontaneous

26. E°cell, G° and K From thermodynamics, we know that, DG° = -RT lnK
We can relate E°cell to free energy for that cell
DG° = -nFE°cell
n = number of moles of e-
So -nFE°cell = -RT lnK and
E°cell = (RT/nF) lnK

27. (Standard state conditions)
Under nonstandard conditions
G = -nFEcell

28. THE NERNST EQUATION Usually concentrations of reactants differ from
one another and also change during the course
of a reaction
As cell reaction proceeds, cell voltage drops so
that E°cell is different from Ecell
E°cell and Ecell are related by the Nernst
Equation
Ecell = E°cell - (RT/nF) lnQ

29. Ecell = E°cell - (RT/nF) lnQ
E = potential under the nonstandard conditions
E° = standard potential
R = gas constant, J/mol.K
T = absolute temperature
n = number of moles of electrons transferred
F = faraday, 96,485 J/V.mol e-
Q = reaction quotient

30. BATTERIES Two type of batteries:
: - Primary batteries cannot be “recharged”
Once all the chemicals are consumed there is
no more chemical reaction
: - Secondary batteries can be regenerated
Most common example is the lead storage
battery used to power automobiles

31. The Lead Storage Battery
Composed of two alternating groups of Pb
plates; one group contains pure lead (anode) and
the other group contains PbO2 (cathode)
The plates are immersed in 40 % sulfuric acid
During discharge
Pb  Pb2+ + 2e- (oxidation)
Pb2+ + SO42-  PbSO4 (precipitation)
Net: Pb + SO42-  PbSO4 + 2e- (anode)

32. At the cathode
PbO2 + 4H+ + 2e-  Pb2+ + 2H2O (reduction)
Pb2+ + SO42-  PbSO4 (precipitation)
Net reaction:
PbO2 + 4H+ + SO e-  PbSO4 + 2H2O
Adding the HR for the two half-cells, gives
Pb + PbO2 + 4H+ + 2SO42-  2PbSO4 + 2H2O
E°cell = V
The battery can be recharged

33. Fuel Cells
These are galvanic cells in which the reactants
are continuously supplied to the cell and the
products are continuously removed
Best known example is the hydrogen-oxygen
fuel cell
Hydrogen is fed into the anode compartment
and oxygen into the cathode compartment

34. Oxygen is reduced at the cathode – porous
carbon doped with metallic catalysts
At the anode hydrogen is oxidized to water
Anode: 2H2(g) + 4OH-(aq)  4H2O(l) + 4e-
Cathode: O2(g) + 2H2O(l) + 4e-  4OH-(aq)
Overall: 2H2(g) + O2(g)  2H2O(g)

35. CORROSION Ordinary corrosion is a redox process in
which metals are oxidized by oxygen in the
presence of moisture
A point of strain on the surface of the metal
acts as an anode
Areas on the metal surface exposed to air
serves as cathodes

36. Anode: Fe(s)  Fe2+(aq) + 2e-
Cathode: O2(g) + 4H+(aq) + 4e-  2H2O(l)
4Fe(s) + O2(g) + 4H+(aq) 
4Fe2+(aq) + 2H2O(l)
2Fe2+(aq) + 4H2O(l)  Fe2O3•H2O(s) + 6H+
Rust
Al also undergo corrosion – initial oxidation
is stopped by a layer of Al2O3

37. Corrosion prevention
1. Plating a metal with a thin layer of a less easily
oxidized metal
2. Allow a protective film to form naturally on the
surface of the metal
3. Galvanizing – coating the metal with zinc
4. Cathodic protection – connecting the metal to
a “sacrificial anode”

38. ELECTROLYTIC CELLS Cells in which an electric current causes a
nonspontaneous reaction to occur – one common
process is called electrolysis
In electrolytic cells the anode is the positive
electrode and the cathode is the negative
electrode
Still : Anode = oxidation; cathode = reduction

39. The Down’ Cell: Electrolysis of molten NaCl
Using graphite inert electrodes the following
observations are made
1. Chlorine, Cl2, is liberated at one electrode
2. Sodium metal forms at the other electrode
Explanation
1. Chlorine is produced at the anode by the
oxidation of Cl- ions

40. 2. Metallic sodium is formed by reducing Na+
ions at the cathode
2Cl  Cl2(g) + 2e- (oxidation, anode HR)
2(Na+ + e-  Na(l) (reduction, cathode HR)
2Na+ + 2Cl-  2Na(l) + Cl2(g) Overall cell rxn.
________________________________________
Electrons used at the cathode are reproduced at
the anode
The reaction is nonspontaneous and electricity
is used to force the reaction to occur

41. Electrolysis of aqueous sodium chloride
In an EChem cell containing aqueous NaCl
: - H2 gas is liberated at one electrode
: - Cl2 gas is liberated at the other electrode
: - Solution at the cathode is basic
Rationalization
: - Chloride ions are oxidized at the anode and
H2O is reduced at the cathode

42. 2Cl-  Cl2 + 2e- (oxidation, anode)
2H2O + 2e-  2OH- + H2 (reduction, cathode)
2H2O + 2Cl-  2OH- + H2 + Cl2 Overall
Sodium metal is more active than hydrogen
metal and liberates H2 from solution
The hydroxide ions are responsible for the
basicity around the cathode

43. FARADAY’S LAW States that the amount of substance that
undergoes oxidation or reduction at each
electrode during electrolysis is directly
proportional to the amount of electricity that
passes through the cell
One faraday = the amount of electricity that
reduces or oxidizes 1 equivalent of a substance

44. One equivalent of any substance is the amount
of that substance that supplies or consumes one
mole of electrons
1F = 1mole of electrons
= x 1023 e-
= 96,485 C
C = It
C= charge passed; I = current;
t = time (in seconds)