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Unit 3
Review
Electrochemistry
Historical Terms
• Reduction is the process of producing metals from their compounds.
• Oxidation is the process of a substance reacting with oxygen
(e.g. combustion, corrosion).
Electron Transfer Theory
Consider the single replacement reaction of zinc
with hydrochloric acid:
Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)
A half–reaction represents what is happening to one
reactant in an overall reaction. It shows either a loss
or gain of electrons by a substance.
(zinc loses two electrons
to become a cation)
Zn(s) → Zn2+(aq) + 2 e–
2 H+(aq) + 2 e– → H2(g)
(two hydrogen ions gain an electron
each and form hydrogen gas)
2 AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2 Ag(s)
A gain of electrons is called reduction.
Ag+(aq) + e– → Ag(s)
A loss of electrons is called oxidation.
Cu(s) → Cu2+(aq) + 2 e–
OIL RIG
“Oxidation Is Loss”
“Reduction Is Gain”
LEO the lion says GER
“Loss of Electrons is Oxidation”
“Gain of Electrons is Reduction”
In any redox (reduction–oxidation) reaction, the total number of
electrons gained in a reaction must equal the total number of electrons
lost.
Writing Complex Half-Reactions
Not all redox reactions (transfer of electrons) are spontaneous.
Cu2+(aq) + Zn(s)
Ni2+(aq) + Cu(s)
spontaneous reaction
no spontaneous reaction
Oxidizing and Reducing Agents
The two entities involved in a redox reaction can be thought of as being
in a “tug of war” for electrons.
The reducing agent causes reduction by donating (losing) electrons.
The reducing agent is oxidized.
“Oxidation Is Loss”
The oxidizing agent causes oxidation by removing (gaining) electrons.
The oxidizing agent is reduced.
“Reduction Is Gain”
Oxidation and reduction are processes.
Oxidizing agents and reducing agents are substances.
The oxidizing agent is reduced.
The reducing agent is oxidized.
Development of a Redox Table
Based on experimental observations, we can rank the ability of selected
metal ions to react with selected metals.
These are reduction reactions.
Instead of looking at the reactivity of the ions, we can use the order of
reactivity of the four solid metals.
These are oxidation reactions.
A redox table is a series of reduction reactions, with the strongest
oxidizing agent at the top left and the strongest reducing agent at the
bottom right of the table.
Remember a reduction reaction is a gain of electrons.
OA + n e

RA
“Oxidation Is Loss”
“Reduction Is Gain”
Based on the table of observations, construct a mini-redox table.
SOA
Br2 (aq) + 2 e 
Ag + (aq) + e
I 2 (aq) + 2 e

Cu 2+ (aq) + 2 e
2 Br  (aq)
Ag(s)

2 I (aq)
Cu(s)
SRA
Cu(s) Mg(s) Ag(s) Zn(s)
Cu2+(aq)
x
Mg2+(aq)
x
x
x
Ag+(aq)
x
x
Zn2+(aq)
SOA
x
x
x
Ag + (aq) + e
2+
Cu (aq) + 2 e

x
Ag(s)
Cu(s)
Zn 2+ (aq) + 2 e
Zn(s)
Mg 2+ (aq) + 2 e
Mg(s)
SRA
The Spontaneity Rule
A student performed the following reactions. Construct a table of
relative strengths of oxidizing and reducing agents.
Co(s) + Pd2+(aq) → Co2+(aq) + Pd(s)
Pd(s) + Pt2+(aq) → Pd2+(aq) + Pt(s)
Mg(s) + Co2+(aq) → Mg2+(aq) + Co(s)
ANSWER:
SOA
Pt 2+ (aq) + 2 e
Pd 2+ (aq) + 2 e
Co2+ (aq) + 2 e
Mg 2+ (aq) + 2 e
Pt(s)
Pd(s)
Co(s)
Mg(s)
SRA
Predicting Redox Reactions
1) List all entities present and
label the possible oxidizing and
reducing agents.
2) Label the strongest oxidizing agent using the chart (highest on the
left) and write the equation for its reduction (following the forward
arrow).
3) Label the strongest reducing agent using the chart (lowest on the
right) and write the equation for its oxidation (following the reverse
arrow).
4) Balance the number of electrons lost and gained in the half-reaction
equations by multiplying one or both of the equations by a number.
Then add the two half-reactions to obtain a net ionic equation.
5) Using the spontaneity rule, predict whether the net ionic equation
represents a spontaneous or nonspontaneous redox reaction.
Another Example
Liquid bromine and chlorine gas are added to a solution of copper(II)
sulphate and a copper strip.
Cl2(g) + Cu(s) → 2 Cl–(aq) + Cu2+(aq)
Yet Another Example!
A few drops of mercury are dropped into a solution which is a mixture
of sulphuric acid and potassium permanganate.
2 MnO4–(aq) + 16 H+(aq) + 5 Hg(l) →
2 Mn2+(aq) + 8 H2O(l) + 5 Hg2+(aq)
Disproportionation
A reaction in which a species is both oxidized and reduced is called
disproportionation.
It may occur when a substance can act
as both an oxidizing agent and as a
reducing agent.
Predicting Redox Reactions by Constructing Half-Reactions
More complex reactions cannot be explained using simple redox theory.
Chemists have developed a method of “electron bookkeeping” to
describe the redox of molecules and complex ions.
The oxidation state of an atom in an entity is defined as the apparent
net electric charge that it would have if electron pairs in covalent bonds
belonged entirely to the more electronegative atom.
e.g. a water molecule
• oxygen has an oxidation state of 2–
• hydrogen has an oxidation state of 1+
An oxidation number is a positive or negative number corresponding
to the oxidation state assigned to an atom in a covalently bonded entity,
based on arbitrary rules.
The sum of the oxidation numbers for a compound (neutral) is zero.
The sum of the oxidation numbers in a polyatomic ion is equal to
the charge on the ion.
Oxidation Numbers and Redox Reactions
An increase in oxidation number is an oxidation.
A decrease in oxidation number is a reduction.
If the oxidation numbers do not change, then a redox reaction has not
occurred (i.e. no transfer of electrons).
Balancing Redox Equations Using Oxidation Numbers
Oxidation numbers and half-reaction equations can be used to balance
any redox equation.
The total increase in oxidation number for a particular atom/ion must
equal the total decrease in oxidation number of another atom/ion.
In a titration, one reagent (the titrant) is slowly
added to another (the sample) until an abrupt change
in solution property (the endpoint) occurs.
In redox titrations, the titrant is
always a strong oxidizing or
reducing agent.
Two oxidizing agents commonly used are acidic
solutions of permanganate ions or dichromate ions.
MnO4–(aq) + 8 H+(aq) + 5 e– → Mn2+(aq) + 4 H2O(l)
purple-pink
colourless
Cr2O72–(aq) + 14 H+(aq) + 6 e– → 2 Cr3+(aq) + 7 H2O(l)
orange
green
Cells and Batteries
Zn
Alessandro Volta
(1745-1827)
Cu
salt water
Volta’s first battery
An electric cell is a device that continuously converts chemical
potential energy to electrical energy.
A battery is a group of two or more electric cells connected to each
other in series.
Volta’s revised battery design (a voltaic pile)
This revised design produces
more electric current for a longer
period of time.
The strips of copper and zinc are replaced
with flat sheets and the jars of brine are
replaced with brine-soaked paper.
Basic Cell Design and Properties
Each cell is composed of two electrodes
(solid electrical conductors) and one
electrolyte (aqueous electrical conductor).
In an electric cell or battery, the
cathode is the positive electrode and
the anode is the negative electrode.
Electrons flow through the wire from the
anode to the cathode.
Consumer, Commercial and Industrial Cells
A primary cell is a cell that cannot be recharged. The reactions are
irreversible and the chemicals become depleted.
The zinc chloride cell is
referred to as a dry cell
because this design was the
first to use a sealed container.
zinc chloride dry cell
2 MnO2(s) + 2 NH4+(aq) + 2 e– → Mn2O3(s) + 2 NH3(aq) + H2O(l)
Zn(s) → Zn2+(aq) + 2 e–
Secondary Cells
A secondary cell can be recharged using electricity to reverse the
chemical reaction that occurs when electricity is produced by the cell.
lead-acid
car battery
PbO2(s) + 4 H+(aq) + SO42–(aq) + 2 e– → PbSO4(s) + 2 H2O(l)
Pb(s) + SO42–(aq) → PbSO4(s) + 2 e–
Fuel Cells
A fuel cell produces electricity by the
reaction of a fuel that is continuously
supplied to keep the cell operating.
hydrogen-oxygen fuel cell
Fuel Cells – Aluminium-Air Cell
Air is pumped into the cell and
oxygen is reduced at the cathode
while aluminum is oxidized at
the anode.
This type of cell has been developed
for possible use in electric cars.
The first electric cells were more for practical purpose of generating
electrical current than scientific study.
To better understand the science behind
the operation of a cell, chemists use a
cell that has two electrodes and their
electrolytes separated.
A porous boundary allows ions to be exchanged between the two
electrolytes but prevents them from mixing.
A voltaic cell is an arrangement of two half-cells separated by a porous
boundary, especially suitable for scientific study.
Zn(s) │ Zn(NO3)2(aq) ║ Cu(NO3)2(aq) │ Cu(s)
│ = phase boundary
║ = physical boundary
Each half-cell consists of an electrode and an electrolyte.
A Theoretical Description of a Voltaic Cell
OA
S OA
Cu(s) │ Cu2+(aq) ║ Ag+(aq) │ Ag(s)
S RA
RA
The strongest oxidizing agent present in the cell
always undergoes a reduction at the cathode.
At the cathode:
• Ag+(aq) + e– → Ag(s)
• electrons enter the cathode through the wire
• silver ions are removed from solution and converted to silver metal
(electrode increases in size)
• to maintain electrical neutrality, cations move through the salt bridge
to the cathode
OA
S OA
Cu(s) │ Cu2+(aq) ║ Ag+(aq) │ Ag(s)
S RA
RA
The strongest reducing agent present in the cell
always undergoes an oxidation at the anode.
At the anode:
• Cu(s) → Cu2+(aq) + 2 e–
• electrons leave the anode through the wire
• copper metal is converted to copper ions in solution
(electrode decreases in size and electrolyte turns more blue)
• to maintain electrical neutrality, anions move through the salt bridge
to the anode
reduction at the cathode:
oxidation at the anode:
net ionic equation:
2 [ Ag+(aq) + e–→ Ag(s) ]
Cu(s) → Cu2+(aq) + 2 e–
Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s)
SUMMARY:
Voltaic Cells with Inert Electrodes
Inert electrodes are used in
voltaic cells where the
oxidizing or reducing agent
is not a solid metal.
They provide a location to
connect a wire and a surface
on which a half-reaction
can occur.
Standard Cells and Cell Potentials
A standard cell is a voltaic cell in which each half-cell contains all
entities shown in the half-reaction equation at SATP, with a
concentration of 1.0 mol/L for the aqueous solutions.
The standard cell potential, E°cell, is the maximum electric potential
difference (voltage) of the cell operating under standard conditions.
E°cell = E°r – E°r
cathode
anode
*must memorize!
Remember, a voltaic cell is a combination of two half-cells
(anode and cathode).
The reduction half-reactions that make up redox table in the data book
are listed in order of decreasing standard reduction potential, E°r.
(standard conditions: 1.0 mol/L solutions at SATP)
The value of the reduction potential, E°r, represents the ability of a
half-cell to attract electrons (i.e. the strength of the oxidizing agent).
The half-cell with the higher reduction potential will attract electrons
from the half-cell with the lower reduction potential.
Standard Hydrogen Half-Cell
It is impossible to determine experimentally the reduction potential of a
single half-cell, because you need two half-cells for a redox reaction to
occur.
The standard hydrogen half-cell has been
chosen as a reference cell, to which all
other reduction potential half cells are
measured relative to.
2 H+(aq) + 2 e– → H2(g)
V
E°r = 0.00
Note that the choice of this half-reaction
as “the zero” completely arbitrary.
Measuring Standard Reduction Potentials
To measure the reduction potential of a half-cell, construct a standard
cell with the hydrogen reference half-cell.
The cell potential, E°cell, is measured with a voltmeter.
If the cell potential is positive, then the positive terminal on the
voltmeter is connected to the cathode and the oxidizing agent at the
cathode is stronger than hydrogen ions.
A positive cell potential (E°cell > 0) indicates that the net
reaction is spontaneous – a requirement for all voltaic cells.
To analyze a standard cell, you
must determine which electrode is
the cathode and which is the anode.
E°cell = E°r –
cathode
anode
E°r
Corrosion
Corrosion is an electrochemical process in which a
metal is oxidized by substances in the environment,
returning the metal to an ore-like state.
In general, any metal appearing below the various
oxygen half-reactions in a redox table will be oxidized
in our environment.
Rusting of Iron
cathode:
anode:
net:
O2(g) + 2 H2O(l) + 4 e– → 4 OH – (aq)
2 [ Fe(s) → Fe2+(aq) + 2 e– ]
2 Fe(s) + O2(g) + 2 H2O(l) → Fe(OH)2(s)
The iron(II) hydroxide precipitate is further oxidized to eventually
form rust: Fe2O3·xH2O(s).
Rusting of Iron
The rusting of iron requires the presence of oxygen and water and is
accelerated by the presence of acidic solutions, electrolytes,
mechanical stresses, and contact with less active metals.
Corrosion Prevention
Protective Coatings
Paint and other metals, such as tin and zinc
are most commonly used.
Tin adheres well to the surface of iron and the
outer surface of the tin coating has a thin, strongly
adhering layer of tin oxide that protects the tin.
Iron has been galvanized when it has
been coated with a layer of zinc.
Zinc is a stronger reducing agent than
iron, thus more easily oxidized.
Fe2+(aq) + 2 e–→ Fe(s)
E°r = –0.45 V
Zn2+(aq) + 2 e–→ Zn(s)
E°r = –0.76 V
Zinc is oxidized instead of the iron it is protecting.
Cathodic Protection
Recall that oxidation (loss of electrons) takes place at the anode of a
voltaic cell.
Cathodic protection is when the iron is forced to become the cathode
by supplying the iron with electrons
Impressed Current Cathodic Protection
An impressed current is an electric
current forced to flow toward an iron
object by an external potential
difference which is provided by a
constant power supply.
This is commonly used for pipelines.
Sacrificial Anode
A sacrificial anode is a metal that is more
easily oxidized than iron and connected to
the iron object to be protected.
The water pipe is turned into the cathode
and magnesium is used as the sacrificial
anode.
Recall, in a voltaic (or electrochemical) cell:
• the strongest oxidizing agent (SOA) is always above the strongest
reducing agent (SRA) in a redox table
• the cell potential, E°cell, is positive
• electrons flow spontaneously from
the anode (–) to the cathode (+).
Electrochemical cells are
made up of two half-cells.
In an electrolytic cell:
• the strongest oxidizing agent (SOA) is always below the strongest
reducing agent (SRA) in a redox table
• the cell potential, E°cell, is negative
• electrons are pulled from the anode (+) to
the cathode (–) by a battery or power supply
• both electrodes are in the same
electrolyte
Electrolysis is the process of supplying
electrical energy to force a nonspontaneous
reaction.
In an electrolytic cell, the chemical reaction is the reverse of that of a
spontaneous cell. During recharging, a secondary cell is functioning
as an electrolytic cell.
The Potassium Iodide Electrolytic Cell: A Synthesis
OA
K+(aq)
I–(aq)
S RA
SOA
H2O(l)
RA
2 H2O(l) + 2 e– → H2(g) + 2 OH–(aq)
E° = –0.83 V
2 I–(aq) → I2(s) + 2 e–
E° = +0.54 V
2 H2O(l) + 2 I–(aq) → H2(g) + 2 OH–(aq) + I2(s)
E°cell = –1.37 V
Evaluation of Predicted Reactions – The Chloride Anomaly
The data table indicates that water is a stronger reducing agent than
chloride ions.
For the electrolysis of solutions containing the chloride ion
(e.g. brine), the chloride ion acts as the strongest reducing
agent, even though the table indicates that it is not.
Science and Technology of Electrolysis
Production of Elements
Most elements occur naturally combined with other elements in
compounds.
The production of active metals (strong reducing
agents) from their minerals typically involves the
electrolysis of molten compounds of the metal.
Sir Humphry Davy
(1778 – 1829)
Many active metals are weaker oxidizing agents than water, so if
aqueous solutions are electrolyzed, water will be reduced instead of
the solid metal you want to produce.
In molten-salt electrolysis, metal cations move to the cathode and are
reduced to metals, and nonmetal anions move to the anode and are
oxidized to nonmetals.
Production of Aluminium
The electrolysis of aluminium
ore is problematic because of the
high melting point of Al2O3(s).
Aluminium oxide dissolves in a
molten mineral called cryolite
(Na3AlF6(l)), which acts as an
inert solvent for the electrolysis
of Al2O3(s).
Overall cell reaction: 2 Al2O3(s) → 4 Al(s) + 3 O2(g)
The Chlor-Alkali Process
The most important nonmetal produced by
electrolysis is chlorine.
The chlor-alkali process is the
electrolysis of brine (NaCl(aq))
to produce chlorine, hydrogen,
and sodium hydroxide.
Remember, the chloride ions (Cl–(aq)) act
as the reducing agent instead of water: the
chloride anomaly.
Refining of Metals
Electrorefining is the process of using
an electrolytic cell to obtain high-grade
metals at the cathode from an impure
metal at the anode.
Electrowinning is the process of using
an electrolytic cell to reduce metal
cations from a molten or aqueous
electrolyte at the cathode.
Electroplating
Electroplating is the deposition of a
metallic coating onto an object by
putting a negative charge onto the
object and immersing it into a solution
which contains a salt of the metal to
be deposited.
The metallic ions of the salt carry
a positive charge and are attracted
to the object.
When they reach it, the negatively
charged object provides the
electrons to reduce the positivelycharged ions to metallic form.
0:48
Oxidation and reduction half-reactions allow us to establish mole ratios
for ions, elements, and electrons.
Two moles of electrons are
required to reduce one mole
2+
–
e.g. Zn (aq) + 2 e → Zn(s)
of zinc ions, producing one
mole of zinc metal.
Q = charge (C)
Q = It
I = current (A)
t = time (s)
The unit of electrical charge
is the coulomb (C).
The unit of electrical current
is the ampere (A).
One coulomb (C) is the quantity of charge transferred by a current of
one ampere (A) during a time of one second.
Faraday’s Law
In an electrolytic cell, the mass of an element produced
or consumed at an electrode is directly proportional to
the time the cell operates at a constant current.
The molar charge of electrons, or the Faraday
constant, F:
C
F  9.65  10
mol e
4
# e- = I x t
F
(data book page 3)
Michael Faraday
(1791 – 1867)
Half-Cell Calculations
Separate calculations are carried out for each electrode, although the same charge, and, therefore, the
same amount of electrons passes through each electrode in a cell or a group of cells in series.