Chapt. 20 – Electrochemistry
19.1 Oxidation and Reduction (review)
19.2 Balancing Redox Equations (intro
only on half reactions)
20.1 Voltaic Cells
20.2 Batteries
20.3 Electrolysis
Section 19.1 Oxidation and Reduction
Oxidation and reduction are
complementary—as an atom is
oxidized, another atom is reduced.
• Describe the processes of oxidation and reduction.
• Identify oxidizing and reducing agents.
• Interpret redox reactions in terms of change in
oxidation state.
Section 19.1 Oxidation and Reduction
Key Concepts
• Oxidation-reduction reactions involve the transfer of
electrons from one atom to another.
• When an atom or ion is reduced, its oxidation
number is lowered. When an atom or ion is oxidized,
its oxidation number is raised.
Section 19.2 Balancing Redox Reactions
Redox equations are balanced when the total
increase in oxidation numbers equals the total
decrease in oxidation numbers of the atoms
involved in the reaction.
• Relate changes in oxidation number to the transfer of
electrons.
Section 19.2 Balancing Redox Reactions
Key Concepts
• The oxidation-number method is based on the number
of electrons transferred from atoms equaling the
number of electrons accepted by other atoms.
• A half-reaction is one of the two parts of a redox
reaction.
Section 20.1 Voltaic Cells
In voltaic cells, oxidation takes place at
the anode, yielding electrons that flow
to the cathode, where reduction occurs.
• Describe a way to obtain electrical energy from a
redox reaction.
• Identify the parts of a voltaic cell, and explain how
each part operates.
• Calculate cell potentials, and determine the
spontaneity of redox reactions.
Section 20.1 Voltaic Cells
Key Concepts
• In a voltaic cell, oxidation and reduction take place at
electrodes separated from each other.
• The standard potential of a half-cell reaction is its voltage when
paired with a standard hydrogen electrode under standard
conditions.
• The reduction potential of a half-cell is negative if it undergoes
oxidation when connected to a standard hydrogen electrode.
The reduction potential of a half-cell is positive if it undergoes
reduction when connected to a standard hydrogen electrode.
• The standard potential of a voltaic cell is the difference between
the standard reduction potentials of the half-cell reactions.
Redox Reactions
Practical /everyday examples:
Corrosion of iron (rust formation)
Forest fire
combustion
Charcoal grill
Natural gas burning
Batteries
Production of Al metal from Al2O3 (alumina)
Metabolic processes
Redox Reactions
GER
LEO:
• Lose Electrons
Oxidation
LEO
LEO says GER
GER:
• Gain Electrons
Reduction
Redox Reactions
Single replacement – zinc in acid
Zn(s) + 2 H+(aq)
0
+1
→
Zn2+ (aq) + H2(g)
+2
0
Zn (s) oxidized
Zn = reducing agent
H+(aq) reduced
H+(aq) = oxidizing agent
Summary of
Terminology for
ReductionOxidation
(Redox)
Reactions
Half-Reactions
Species – any kind of chemical unit involved
in a chemical reaction
NH3(g) + H2O(l)  NH4+(aq) + OH-(aq)
4 species in above: 2 molecules & 2 ions
Redox reaction occurs when species that
can give up electrons comes in contact with
species that can accept them
Half-reaction: one of 2 parts of redox
equation (oxidation or reduction half)
Half-Reactions
Redox equation
2Fe(s) + 3Cl2(g)  2FeCl3(s)
Half-reactions:
Oxidation: Fe  Fe3+ + 3eReduction: Cl2 + 2e-  2Cl-
Redox Reaction Example
On right—Zn metal
is dipped in Cu2+
solution (blue)
After a bit, blue
solution is lighter,
and metal strip is
covered with Cu
metal
Redox Reaction Example
Viewed as overall single replacement
reaction (net ionic):
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Viewed as separate half reactions:
Zn(s) → Zn2+(aq) + 2e- (oxidation)
Cu2+(aq) + 2e- → Cu(s) (reduction)
Moving electrons from one species to
another
Reductionist View of Redox Reactions
Movement of electrons from a source to a
recipient
Half Reactions – Oxidation of Iron
Oxidation half reaction: Fe → Fe3+ + 3e–
Overall Rxn (unbal.)
Reduction Half Rxn
Fe + O2 → Fe2O3
O2 + 4e– → 2O2 –
Fe + Cl2 → FeCl3
Cl2 + 2e– → 2Cl–
Fe + HBr → H2 + FeBr3
2H+ + 4e– → H2
Fe + AgNO3 →
Ag + Fe(NO3)3
Fe + CuSO4 →
Cu + Fe2(SO4)3
Ag + e– → Ag
Cu2+ + 2e– → Cu
Terminology
Simple electrochemical cells are
referred to as galvanic cells (after
Galvani) or voltaic cells (after Volta)
Our textbook uses term voltaic
Voltaic Cells
What happens if Zn metal immersed in
CuSO4(aq) solution?
Spontaneous* redox reaction occurs:
Zn(s) +

+ Cu(s)
* doesn’t need to be driven by outside
source of energy (DG for process < 0)
+
2
Zn(s) oxidized and Cu (aq) reduced
+
2
Cu (aq)
+
2
Zn (aq)
Half Reactions → Half Cells
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
For this example, redox half reactions occur
in same place (species in direct contact)
Can separate them as half cells by
constructing electrochemical cell (indirect)
Electrochemical Cell
An apparatus that uses a redox
reaction to produce electrical energy or
uses electrical energy to cause a
chemical reaction
Voltaic (Galvanic) Cells
Device in which spontaneous redox
reaction occurs as electrons are
transferred from reductant to oxidant
through an external circuit
Used to perform electrical work using
energy released during spontaneous
redox reaction
Voltaic Cell: Half-Cells
Two half reactions occur in separate
compartments called half-cells
• 1 half-cell contains oxidation half reaction
• 1 half-cell contains reduction half reaction
Each half cell
contains:
• metal electrode
• electrolyte solution
of ion of electrode
1M
Zn2+
1M
Cu2+
Voltaic Cell: Half-Cells
Connect half-cells
Redox happens electrons can now
transfer from 1
electrode to another
e-
e-
Zn
-
e
e-
e-
Cu
e-
ee-
eZn+2
Cu+2
Process stops almost immediately –
buildup of excess negative charge on
Cu electrode prevents further transfer
Electron flow
Ion flow keeps the charge
neutral
Salt bridge contains strong electrolyte
Voltaic Cell: Complete Cell
Add salt bridge – allows ion transport
Bridge has soluble salt
(e.g., KCl) in agar gel
(like jello) with ion
permeable plugs
Ion movement
neutralizes charge
created by electron
movement
e-
e-
KCl Salt Bridge
Cl-   K+
ee-
Cl-
e-
K+
e-
Cl-
K+
Porous Disk Also Allows Ion Flow
Voltaic Cell Using Porous Barrier
Daniell cell: Cu & Zn
electrodes dipping into
solutions of copper(II)
sulfate and zinc
sulfate, respectively
Solutions make contact
through porous pot,
which allows ions to
pass through to
complete circuit
Anodes and Cathodes
Anode
Electrode at which oxidation occurs
• located in oxidation half-cell
• the “negative” electrode
• electrons are released here
• anions move toward anode
Anodes and Cathodes
Cathode
Electrode at which reduction occurs
• located in reduction half-cell
• the “positive” electrode
• electrons move toward here
• cations move toward anode
General Voltaic Cell
Anode
half-cell
Cathode
half-cell
Oxidation
occurs
here
Reduction
occurs
here
Voltaic Cell (with nitrates)
Zn(s) + Cu2+(aq)
Oxidation
half cell
Zn2+(aq) + Cu(s)
Reduction
half cell
Anodes and Cathodes
For following reaction:
Zn(s) + Ni2+(aq) 
Zn2+(aq) + Ni(s)
• Which metal will be the anode?
Oxidation occurs at anode: Zn (s)
• Which metal will be the cathode?
Reduction occurs at cathode: Ni (s)
Spontaneity and Potential Energy
Redox reactions occurring in voltaic cell
are spontaneous
Why do electrons flow spontaneously
from one electrode to other?
• Flow spontaneously due to difference in
potential energy between anode and
cathode
Current, Voltage (Cell Potential), and
Electrical Potential Energy
Electric potential energy: measure of
amount of current (flow of charge) that can
be generated from voltaic cell to do work
Current can flow between two points only
when a difference in electric potential
energy exists between the two points
Volt: unit used to measure cell potential—
driving force from difference in electric
potential energy between two electrodes
Electron Flow and Potential Energy
Anode
• higher
potential
energy
Cathode
• lower
potential
energy
Height is an analogy for voltage
High Voltage
Low Voltage
Current = charge passing/unit time
Low Amps
High Amps
Voltaic Cells – Cell Potential
Difference in electrical potential
between anode and cathode called
electromotive force (emf)
• Also known as cell potential or cell
voltage
• measured in volts (V)
• Indication of energy available to move
electrons from anode to cathode
Voltaic Cells – Standard emf
Standard cell potential (emf) (Eocell)
Cell potential measured under standard
conditions
• 25oC
• 1M concentrations of reactants and
products in solution
• 1 atm pressure for gases
Voltaic Cells -
o
E
cell
Eocell depends on half-cells or halfreactions present
Standard potentials have been
assigned to each individual half-cell –
cannot be directly determined
By convention, standard reduction
potential (Eored) for each half cell is
used
Voltaic Cells -
o
E
cell
Must measure reduction potential of
half cell against something – a
reference
Use standard hydrogen electrode
Platinum electrode in 1 M HCl, 25oC, 1
atm H2(g)
Differs from previous half cells – metal
not involved in redox process
Standard Hydrogen Electrode
Redox process for SHE:
2H+ (aq, 1M) + 2 e-
 H2 (g, 1 atm)
Reduction potential (at 25oC) of this
half-cell assigned a potential of exactly
0.000 V
A Hydrogen
Electrode
= SHE if
[H+] = 1 M
(from HCl),
p(H2(g)) = 1 atm,
T = 25oC
Cell Potentials – Cu vs SHE
e- flow
Cu =
cathode
(reduction)
Eocell =
+0.342 V
e
e
Cu2+
1 M Cu2+
H+
1 M H+
Cell Potentials – Zn vs SHE
e- flow
Zn =
anode
(oxidation)
Eocell =
0.762 V
Cu2+
e- flow
H+
e
e
Zn2+
H+
1 M Zn2+
1 M H+
o
E
red
Values
Eored tabulated table 25.1, p 712
Ordered from most negative (Li+ strong reducing agent) to most positive
(F2 – strong oxidizing agent)
Includes many reactions that don’t
involve metals
O2 + 2H+ +2e-  H2O2
For these half-cells, unreactive metal
used as electrode (e.g., platinum)
o
E
red
Values (table 20.1)
Table of Standard Reduction Potentials
Eo for reaction as written
The more positive Eo,
the greater tendency for
substance to be reduced
Eo
Sign of
changes when
reaction reversed
Changing stoichiometric
coefficients of half-cell
reaction does not
change value of Eo
Eo (V)
Cu2+ + 2e-
Cu
+0.34
2 H+ + 2e-
H2
0.00
Zn2+ + 2e-
Zn
-0.76
Greater reducing tendency
Sign of
o
E
red
and Spontaneity
As Eored becomes increasingly positive,
driving force for reduction increases
F2(g) + 2e-  2 F-(aq) Eored = +2.87 V
Ag+(aq) + e-  Ag(s)
Eored = +0.80 V
Fluorine (most EN element) wants to be
reduced more than silver ion
Sign of
o
E
red
and Spontaneity
As Eored becomes increasingly
negative, driving force for oxidation
increases
Li+ (aq) + e-  Li (s) Eored = -3.05 V
Negative reduction potential indicates
that oxidation half-reaction is
spontaneous
Li (s)  Li+ (aq) + e-
Standard Reduction Potentials
Given following potentials, which metals will
be most easily oxidized?
 Ag(s)
Eored = +0.80 V
Zn2+(aq) + 2 e-  Zn(s)
Eored = - 0.76 V
 Na(s)
Eored = - 2.71 V
Ag+(aq) + eNa+(aq) + e-
?
Na(s) most easily oxidized
Potentials are quantitative version of “metal
activity series” previously used
Cell Potential – Anodes & Cathodes
Make voltaic cell by combining two half-cells
and calculate standard cell potential, Eocell
Cell must have both anodic process (anode)
and cathodic process (cathode)
Look at standard reduction potential, Eored
for each half-cell: cell with more positive
value occurs as written (as a reduction =
cathode)
Other half-cell will occur in reverse (as an
oxidation = anode)
Standard cell potential,
o
E
cell
Once anode and cathode have been
identified, calculate Eocell from
Eocell = Eored (cathode) – Eored (anode)
reduction
oxidation
Book uses equation
Eocell = Eoreduction – Eooxidation
Misleading - second term is still really a
reduction potential (from table)
Example: Eocell for Cu/Zn
Step 1 – Identify anode and cathode
Cu2+(aq) + 2 e-  Cu (s) Eored = +0.342 V
Zn2+(aq) + 2 e-  Zn(s) Eored = - 0.762 V
Potential for Cu more positive
Cu will be cathode  Zn will be anode
Step 2 – Calculate standard cell potential
Eocell = Eored (cathode) – Eored (anode)
Eocell = 0.342 V - (-0.762 V)
Eocell = 1.104 V
Overall Cell Potential – Cu/Zn Cell
Zn│Zn+2║Cu+2│Cu
anode
cathode
Cell notation
E0Cu = +0.342 V
E0SHE = 0.000 V
E0cell = +1.104 V
E0Zn =  0.762 V
Voltaic Cells – Line Notation
Anode components listed on left
Cathode components listed on right
Anode and cathode have reactants on
left, products on right
Separate half cells with double vertical
lines: ll
Indicate phase difference with single
vertical line: l
Line Notation & Overall Reaction
anode|anode solution║cathode solution|cathode
Al(s) | Al3+ (1.00 M) ║Cu2+ (1.00 M) | Cu(s)
Anode: Al(s)  Al3+(aq) + 3e- Eored = -1.662 V
Cathode: Cu2+(aq)  Cu(s) + 2e- Eored = 0.342 V
Note: Cell notation does not denote balance
To get balanced overall reaction, multiply anodic
×2 and cathodic ×3 to give
2Al(s) + 3Cu2+(aq)  3Cu(s) + 2Al3+(aq)
Eocell = Eored(cathode) – Eored(anode)
Eocell = 0.342 V – (-1.662 V) = 2.004 V
Practice
Using Cell (Line) Notation
Problems 34, 39(a-d), 40(a-c) page
736
Voltaic Cells
Given following reduction half-cells, identify
anode, cell line notation, balanced reaction
for cell, and Eocell
Al3+(aq) + 3 e-  Al(s) Eored = -1.66 V
Fe2+(aq) + 2 e-  Fe(s) Eored = -0.440 V
More positive half-cell = Fe  cathode
So Al is anode: Al (s)  Al3+ (aq) + 3 eanode|anode solution║cathode solution|cathode
Al|Al+3║Fe2+|Fe
Voltaic Cells
Balanced equation:
3+(aq) + 3 eAl(s)

Al
Anode:
×2
2+(aq) + 2 e-  Fe(s)
Fe
Cathode:
×3
2 Al(s) + 3 Fe2+(aq) + 6e-  2 Al3+(aq) + 3 Fe(s) + 6e-
2 Al(s) + 3 Fe2+(aq)  2 Al3+(aq) + 3 Fe(s)
Eocell = Eored (cathode) - Eored (anode)
Eocell = -0.440 V  (-1.66 V) = 1.22 V
Reduction potentials not multiplied by anything
Practice
Calculating standard cell potentials
Problems 1 – 4, page 716
Problem 41(a-c) page 736
Problem 68 page 738
Problems 1(a-c) page 991
Is Proposed Reaction Spontaneous?
Used E0red data to get E0cell; process
guaranteed to give spontaneous overall
reaction (E0cell always > 0)
Can reverse process to determine if a given
redox reaction is spontaneous as written
Step 1 – Write reaction in form of half-cells
Step 2 – Find E0red for each
Step 3 – Calculate E0cell
Step 4 – If E0cell > 0, reaction spontaneous
if < 0 not spont., but reverse rxn is
Practice
Determining spontaneity of redox
reaction
Problems 5 – 9, page 716
Problems 13(a-c) page 717
Problems 67(a-d), page 738
Problems 2(a-c) page 991
Chapt. 20 – Electrochemistry
19.1 Oxidation and Reduction (review)
19.2 Balancing Redox Equations (intro
only on half reactions)
20.1 Voltaic Cells
20.2 Batteries
20.3 Electrolysis
Section 20.2 Batteries
Batteries are voltaic cells that use
spontaneous reactions to provide
energy for a variety of purposes.
• Describe the structure, composition, and operation of
the typical carbon-zinc dry-cell battery.
• Distinguish between primary and secondary batteries,
and give two examples of each type.
• Explain the structure and operation of the hydrogenoxygen fuel cell.
• Describe the process of corrosion of iron and methods
to prevent corrosion.
Section 20.2 Batteries
Key Concepts
• Primary batteries can be used only once; secondary
batteries can be recharged.
• When a battery is recharged, electric energy supplied
to the battery reverses the direction of the battery’s
spontaneous reaction.
• Fuel cells are batteries in which the substance
oxidized is a fuel from an external source.
• Methods of preventing corrosion are painting, coating
(plating) with another metal, or using a sacrificial
anode.
Battery
A battery is a:
http://www.powerstream.com/BatteryFAQ.html
• voltaic (galvanic) cell
• or group of voltaic cells connected in
series

cell potentials of individual cells add up to
give total battery cell potential
• source of direct current (DC)
Batteries – Primary vs Secondary
Primary – cannot be recharged
because one or both half-reactions not
reversible
• can explode if recharge attempted
Secondary – can be recharged by
reversing flow of current, regenerating
reactants
• AKA storage batteries
Batteries – Wet Cell vs Dry Cell
Wet cell – conventional electrolyte
“Dry” Cell – typically a moist paste
Alessandro Volta
Showed that If put
electrolyte–soaked
cloth in between 2
different metals,
(e.g., Cu & Zn),
current would flow
Types of Primary Cell Batteries
Alkaline
Aluminium (aluminum-air)
Lithium (not Lithium-ion)
Oxyride (Oxy Nickel Hydroxide (NiOOH)
Silver-oxide
Zinc-air
Zinc-carbon
Zinc-Carbon Dry Cell
Wet version 1866, “dry” version 1887
Anode: Zn
Cathode : carbon rod (inactive) in contact
with moist, acidic paste of solid MnO2, solid
NH4Cl, and powdered graphite
Produces 1.5 V
Inexpensive, safe
Short shelf life due to reaction of Zn with
acidic electrolyte
Can generate NH3(g) at high current drains
Standard
Zinc-Carbon
Dry Cell
Battery
Simplified overall
reaction
2NH4+(aq) + 2 MnO2(s) + Zn(s) 
Zn+2(aq) + Mn2O3(s) +2 NH3(aq) + H2O(l)
Zinc-Carbon: Half & Other Reactions
Zn(s)  Zn2+(aq) = 2e-
(anode)
2MnO2 (s) + 2NH4+(aq) + 2e-  Mn2O3(s)
+ 2NH3(aq) + H2O (l)
(cathode)
Zn2+(aq) + 2NH3 (aq) + 2Cl-(aq) 
Zn(NH3)2Cl2(s) (complexation – no redox)
Overall:
2MnO2 (s) + 2NH4Cl(aq) + Zn(s) 
Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s)
Alkaline Battery
Anode: Zn powder (not solid)
Cathode: moist paste of solid MnO2, KOH or
NaOH, and carbon (no NH4Cl as in normal
carbon-zinc) with current collector (various)
Lasts longer than standard carbon-zinc:
zinc doesn’t corrode as fast under basic
conditions
Has less voltage drop than carbon-zinc
More expensive; prone to leaking KOH
Alkaline Battery
Zn-KOH anode
paste
Brass current
collector
KOH
electrolyte
Steel Case
MnO2, cathode (current
mixture
collector)
Alkaline Battery
E°Cell = 1.5 V
Published half reactions vary [book uses
first anode and first cathode reactions]:
Anode: Zn(s) + 2OH-(aq)  ZnO(s) + H2O(l) + 2eAnode: Zn(s) + 2OH-(aq)  Zn(OH)2(s) + 2eCathode (+4  +2):
MnO2 (s) + 2H2O(l) + 2e-  Mn(OH)2 (s) + 2OH-(aq)
Cathode(+4  +3):
2MnO2 (s) + H2O(l) + 2e-  Mn2O3 (s) + 2OH-(aq)
Lithium Batteries
Wide variety of types of cathodes and
electrolytes
All types have lithium metal as anode [book
does not properly distinguish between
lithium & lithium-ion (no solid Li) batteries!]
Low metal density makes lightweight battery
Low standard reduction potential gives
higher output voltage (2.3 V higher than
with cell utilizing Zn anode) – typical 3 V out
Long lasting
Lithium Batteries
Most common: MnO2 cathode, salt of lithium
(LiClO4) dissolved in organic solvent
(propylene carbonate and dimethoxy
ethane) as electrolyte
For medical use: Li-I2, Iodine cathode, solid
electrolyte (charge transfer complex, e.g.,
poly-2-vinylpyridine)
Secondary Batteries
Secondary – can be recharged by
reversing flow of current, regenerating
reactants
• AKA storage batteries
Types of Secondary Cell Batteries
Lead-acid (automobile)
Lithium
Lithium-ion
Lithium-ion polymer
Nickel-cadmium (NiCad)
Nickel metal hydride (NiMH)
Molten salt
Lead-Acid (Storage) Battery
Most common secondary battery
Lead Acid Battery
Can function for several years over
temperature range from -30oF to 120oF
12 V battery (six cells)
anode = Pb
cathode = Pb coated with PbO2
electrolyte solution = H2SO4 solution
Lead storage battery
Lead Acid Battery (+2.041 V)
Pb(s) + SO42-(aq) → PbSO4(s) + 2 e- +1.685 V
PbO2(s) + SO42-(aq) + 4H+(aq) + 2 e- → PbSO4(s)
+ 2 H2O(l)
+0.356 V
Pb + PbO2 + 2SO42- + 4H+ → 2 PbSO4 + 2 H2O
PbSO4 (from both electrodes) adheres
During discharge, H2SO4 consumed at both
electrodes and H2O produced
Battery condition can be determined by
measuring electrolyte density – drops with
drop in acid concentration
Lead Acid Battery
Can be recharged because PbSO4(s)
products adhere to electrodes - alternator
can force current through battery in
opposite direction and reverse reactions
Even though battery can be recharged,
physical damage from road shock and
chemical side reactions (e.g. electrolysis of
water) eventually cause battery failure
NiCad Battery
1.4 V
Rechargeable (like lead acid
battery, products adhere to
electrodes)
NiCad Cell
Cathode (Ni3+ to Ni2+)
NiOOH(s) + H2O(l) +
e–
–
Ni(OH)2(s) + OH (aq)
Anode (Cd to Cd2+)
–
Cd(s) + 2 OH (aq)
Cd(OH)2(s) + 2 e–
Overall
2 NiOOH(s) + Cd(s)
2 Ni(OH)2(s) + Cd(OH)2(s)
Lithium-Ion Cell
3.6 V
One of most popular types for portable
electronics (Sony commercialized in 1991)
High energy to weight ratio (2x NiCad)
No memory effect
Low self-discharge (5% month) compared to
NiCad (10%) and NiMH (30%) but subject to
aging starting from time of manufacture
Solid, polymeric electrolyte (salt bridge)
solid electrolyte interphase (SEI).
solid lithium-salt electrolytes (LiPF6, LiBF4, or LiClO4) and organic solvent
Lithium-Ion Cell
Anode: carbon (graphite)
Cathode: metal oxide of Li with Co (most),
Mn, or Ni/Co/Mn (NCM) – no metallic Li
Electrolyte: lithium salt in organic solvent
Reaction:
Li+ transported to and from cathode/anode
Transition metal, Co, in LixCoO2 oxidized
from Co3+ to Co4+ during charging, and
reduced from Co4+ to Co3+ during discharge
Fuel Cell
Voltaic cell for which reactants are
continuously supplied
Used in U.S. space program
Based on reaction of hydrogen (and other
fuels such as methane) with oxygen to form
water
2 H2(g) + O2(g) → 2H2O(l)
Same as combustion, but done in way that
electrical energy can be extracted
Fuel Cell
Electrodes:
Hollow chamber of porous carbon walls
Walls of chamber contain catalysts (Pt, Pd)
Electrolyte:
KOH (alkaline cell)
Or
Proton-exchange membrane (PEM) – allow
H+ ions to pass through
Safer & lighter than using liquid electrolyte
Fuel Cell with KOH Electrolyte
Anode: 2H2 (g) + 4OH- (aq) → 4H2O (l) + 4eCathode: O2 (g) + 2H2O (l) + 4e- → 4OH- (aq)
2H2 (g) + O2 (g) → 2H2O (l)
Hydrogen Fuel Cell with PEM
Corrosion
Loss of metal resulting from redox
reaction of metal with substances in
environment
Ordinary rusting requires presence of
both oxygen and water (electrolyte)
Iron surface naturally becomes
inhomogeneous and develops anodic
and cathodic regions
Corrosion Cell
and often
Fe2+  Fe3+ + e-
Corrosion
Oxidation: Fe(s)  Fe2+(aq) + 2 e–
Reduction: O2(g) + 4 H+(aq) + 4 e–  2 H2O(l)
Overall: 2 Fe(s) + O2(g) + 4 H+(aq) 
2Fe2+(aq) + 2 H2O(l)
Corrosion
Can be slow in pure water
Faster in water with dissolved salts,
especially with chloride or sulfate ions
Seawater especially corrosive
Corrosion Prevention
Organic coatings (paint) can be
effective
Problems occur at breaks, pinholes
Alternative: sacrificial anodes
Metal with more negative reduction
potential than iron – Mg, Zn
This material corrodes and iron is
cathode only (stays protected)
Use of Zinc Sacrificial Anode
Sacrificial Mg Anode to Protect
Steel Pipeline
Cathodic Protection of Iron Storage
Tank Using Mg Sacrificial Anode
Corrosion Prevention - Galvanizing
Combines coating idea (barrier) with
sacrificial anode idea
Inexpensive and widely used
Especially important for parts of car
bodies
Intact Zinc Coating (Barrier)
Breached Zinc Coating
(Sacrificial Anode)
Chapt. 20 – Electrochemistry
19.1 Oxidation and Reduction (review)
19.2 Balancing Redox Equations (intro
only on half reactions)
20.1 Voltaic Cells
20.2 Batteries
20.3 Electrolysis
Section 20.3 Electrolysis
In electrolysis, a power source causes
nonspontaneous reactions to occur in
electrochemical cells.
• Describe how it is possible to reverse a spontaneous
redox reaction in an electrochemical cell.
• Compare the reactions involved in the electrolysis of
molten sodium chloride with those in the electrolysis of
brine.
• Discuss the importance of electrolysis in the smelting
and purification of metals.
Section 20.3 Electrolysis
Key Concepts
• In an electrolytic cell, an outside source of power
causes a nonspontaneous redox reaction to occur.
• The electrolysis of molten sodium chloride yields
sodium metal and chlorine gas. The electrolysis of
brine yields hydrogen gas, sodium hydroxide, and
chlorine gas.
• Metals such as copper are purified in an electrolytic
cell.
• Electrolysis is used to electroplate objects and to
produce pure aluminum from its ore.
Electrolysis
Use of electrical energy to bring about a
chemical reaction
Electrolysis cell: electrochemical cell in
which electrolysis occurs
Reverse of normal voltaic (galvanic) cell
process - apply voltage to drive reactions in
reverse (same as charging secondary
battery)
Not possible for some systems (primary
batteries)
Electrolytic Cell Discharge (a)
and Recharge (b)
Note: this particular cell is not really rechargeable
due to migration of Zn2+(aq) away from electrode
e
flow
e flow
Electrolysis Applications
Water – produces H2(g), O2(g)
Down’s cell – molten NaCl to produce
Na(l), Cl2(g)
Chloralkali process – electrolyze brine
to generate NaOH and Cl2(g)
Aluminum production (Hall process)
Electrorefining (copper)
Electroplating
Cell for electrolysis of water
Generates H2(g) & O2(g) in 2:1 ratio
Electrolysis Applications
Down’s cell – molten NaCl to produce
Na(l), Cl2(g)
Electrolysis
of Molten
NaCl –
Down’s Cell
Only practical way to
get Na
Na+(l)  Na(l) + e2Cl-(l) Cl2(g) + 2e2Na+(l) + 2Cl-(l) 
2Na(l) + Cl2(g)
Brine Electrolysis
Chloralkali Production
Brine = concentrated aqueous NaCl
Overall reaction:
2H2O(l) + 2NaCl(aq) 
H2(g) + Cl2(g) + 2NaOH(aq)
All 3 products commercially important
Brine Electrolysis
Producing Aluminum
2 Al2O3 + 3 C  4 Al + 3 CO2
Charles Hall (1863-1914) developed process
Founded Alcoa
Price of Al dropped from $100,000/lb in 1855 to
$2/lb in 1890.
Producing Aluminum
L.T.Héroult discovered same process at
same time; called Hall-Héroult process
Electrolyze Al2O3 (from bauxite) - dissolved
at 1000C in molten cryolite (Na3AlF6)
Cell lined with graphite = cathode:
Al3+ +3e-  Al(l)
Graphite rods are also anode:
2O2- O2(g) + 4e- & 2C(s) + O2(g) CO(g)
Requires huge amount of energy – Al
production uses 4% of US electrical power
Refining of Aluminum
Refining Al – Hall-Héroult Process
Power
Source
Copper Electrorefining
Anodes: slabs of impure Cu
Cathodes: thin sheets of pure Cu
Electrolyte: acidic copper sulfate
Voltage across electrodes designed to
produce only Cu at cathode
Metallic impurities that can oxidize to ions
(e.g., Zn) do not plate out on cathode
Insoluble metal impurities (don’t oxidize)
collected in sludge at bottom of cell
Copper Electrorefining
Impure Cu
Anode
Cu(s) 
Cu2+(aq) + 2eand other
metals more
easily oxidized
than Cu
Pure Cu
Cathode
Cu2+(aq) + 2e Cu(s)
Voltage in
range where
this only can
happen for Cu
Electroplating
Used to produce decorative and/or
protective layer of metal on top of a
second (usually cheaper) metal
Metal being plated is cathode
Source of metal to plate may be anode
or metal ion in solution (Au,Cr)
Electroplating
Electroplating
of silver
End of Chapter