Chapter 18: Electrochemistry

1
Chapter 18 Electrochemistry

• Area concerned with the interconversion of
• chemical and electrical energy

2
Electrochemistry

• 2 types of electrochemical cells
• 1.  Galvanic cells (voltaic cells)
• Spontaneous chemical reaction generates an
  electric current
• 2.  Electrolytic cells
• An electric current drives a non-spontaneous
  reaction
• These two types are the reverse of each other

3
Oxidation a loss of electrons (an increase in
oxidation number Reduction a gain of electrons
(a decrease in oxidation number) Redox overall
reaction, one partner loses electrons, other
partner gains electrons Oxidation and reduction
are half-reactions must be combined to give the
overall reaction Metallic state ionic
state Au Au
e noble Cu Cu2 2e semi-noble Zn Zn
2 2e base Na Na e very base
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There is a natural tendency of substances to gain
or lose electrons Comparing the two metals Cu
and Zn Cu is more likely to be in the metallic
state Zn is more likely to be in the ionic
state Zn CuSO4 ? ZnSO4
Cu net Zn Cu2 ? Zn2
Cu Cu ZnSO4 ? NR
5
Review of Redox

• Overall reaction
• Zn Cu2 ? Zn2 Cu
• ½ ox. rxn Zn ? Zn2 2e-
• ½ red. rxn Cu2 2e- ? Cu
•  
• Cu2 is reduced and it is the oxidizing agent
• Zn is oxidized and it is the reducing agent
• Electrons are transferred directly from Zn to
  Cu2
• 4. Enthalpy of rxn is lost to the surroundings
  as heat

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Same reaction with electrochemical cell
(galvanic) Some of the chemical energy released
by the reaction is converted to electrical
energy which can be used to light a light bulb

Apparatus
7

• Electrodes
• strips of zinc and copper in separate
  compartments
• connected by an electrically conductive wire
• electrons transferred through wire make current
• Anode is negative, oxidation takes place (Zn
  strip)
• Cathode is positive, reduction takes place (Cu
  strip)
• Salt bridge
• U-shaped tube that contains a gel permeated with
  a solution of
• inert electrolyte (will not react)
• Necessary to complete the electrical circuit
•  
• Anode Ox ½ Zn ? Zn2 2e-
• Cathode Red ½ Cu2 2e- ? Cu
• Overall Rxn Zn Cu2 ? Zn2 Cu

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• Why is the salt bridge necessary?
• 1. Completes the electrical circuit
• Without it the anode cell would become
  positively charged as
• Zn2 ions appear the solution in the cathode
  beaker would
• become negatively charged as Cu2 ions are
  removed
• Because of the charge imbalance the elctrode
  reactions would
• quickly stop and electron flow through the wire
  stops
• With the salt bridge the electrical neutrality is
  maintained in
• both beakers by a flow of ions
• Anions- SO42- flow through salt bridge from
  cathode to anode
• Cations- from anode to cathode

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Overview

• Anode Oxidation occurs
• electrons are produced
• anions move towards anode
• negative sign
• Cathode Reduction occurs
• electrons are consumed
• cations move towards
• positive sign
• Electrons move from the anode (where they are
• produced) to the cathode (where they are
  consumed)

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Shorthand Notation for Galvanic Cells

• Example Zn(s) Cu2 ? Zn2 Cu(s)
• Shorthand notation Zn(s) ? Zn2(aq) ?? Cu2(aq)
  ? Cu(s)
• 1. ?, single vertical line represents a phase
  boundary (between a solid electrode and an
  aqueous solution
• 2. ??, double vertical line denotes a salt
  bridge
• 3. The anode half-cell is always on the left of
  the salt bridge, with the solid electrode to the
  far left
• 4. The cathode half-cell is always on the right
  of the salt bridge, with the solid electrode on
  the far right
• 5. The reactants in each half cell are always
  written first, followed by the products
• 6. Electrons move through the external circuit
  from left to right (from anode to cathode)

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Example 1

• Write the shorthand notation for a galvanic cell
  that uses the reaction
•  
• Fe(s) Sn2(aq) ? Fe2(aq) Sn(s)

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Cell Potentials and Free-Energy Changes for Cell
Reactions

• Electrons move through the external circuit from
  the zinc anode to the copper cathode because they
  have lower energy when on copper than on zinc.
  The driving force that pushes the negatively
  charges electrons away from the anode (-
  electrode) and pulls them toward the cathode (
  electrode) is an electrical potential called the
  electromotive force (emf) also known as cell
  potential (E) or the cell voltage. SI units in
  volts (V)

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Definitions Electrical charge unit Coulomb C
1C 1As (Ampere x sec) Faraday constant F
electrical charge of 1 mol of electrons 96,485
C Energy Unit Joule 1 J 1 C x 1 V (charge
x voltage) (Work required to move an electrical
charge of one coulomb Through an electrical
potential difference of one volt.) Gibbs free
energy DG - n F E n number of moles
of electrons n F charge transferred E
cell potential
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Cell Potentials and Free-Energy Changes for Cell
Reactions

• ?G -nFE
• Standard cell potential E0 reactants and
  products are in their
• standard state (1 M concentration, gases at 1
  atm, temp 250C)
• ?G? -nFE?
• Standard free-energy change and standard cell
  potential
• Because both are directly proportional a
  voltmeter can be regarded as a free- energy
  meter
• When a voltmeter measures E?, it is also
  indirectly measuring ?G?

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• The standard cell potential at 25?C is 0.92 V for
  the reaction
•  
• Al(s) Cr3(aq) ? Al3(aq) Cr(s)
•  
• What is the standard free-energy change for this
  reaction at 25?C in kJ?
• ?G? -nFE? - 3 mol e x 96,500 C/mol x 0.92
  J/C V
• DGo - 266.3 kJ
• Is the process spontaneous?
• Which metal is more noble?

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The standard potential of any galvanic cell is
the sum of the standard half-cell potential for
oxidation at the anode standard half-cell
potential for reduction at the cathode
E?cell E?ox E?red However, we cannot
measure the potential of a single electrode We
can only measure a voltage potential
difference between two electrodes with a
voltmeter We need an arbitrary standard
half-cell as a reference point and assign an
arbitrary potential
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Standard hydrogen electrode (SHE) is used as the
arbitrary half-cell Definition the SHE has a
standard potential of 0.00 Volt (sea level)
1 platinum electrode 2 hydrogen blow 3 sulfuric
acid with pH 0
Half-cell reaction H2(g) ? 2 H(aq) 2 e EoH2
0.00 Volt
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• The standard potential of any galvanic cell is
  the sum of the standard half-cell potentials for
  oxidation at the anode and reduction at the
  cathode
• E?cell E?ox E?red

H2 gas is oxidized to H ions at the anode and
Cu2 ions are reduced to copper metal at the
cathode Which element is more noble?
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• Anode (oxidation) H2(g) ? 2 H(aq) 2 e
• Cathode (reduction) Cu2(aq) 2 e ?
  Cu(s)
•  
• Overall H2(g) Cu2(aq) ? 2
  H(aq) Cu(s)
•  
• E?cell E?ox E?red E?H2?H E?Cu2 ? Cu
  
• 0 V .34 V .34 V

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Now we measure the standard cell potential for
Zinc Cathode (reduction) 2 H(aq) 2 e ?
H2(g) Anode (oxidation) Zn(s) ?
Zn2(aq) 2e   Overall 2 H (aq)
Zn(s) ? H2(g) Zn2(s)  Eocell .76 V
Which element is more noble in this case?
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• Standard potential for an oxidation half-reaction
  is the
• negative of the standard reduction potential.
• Cu2 2 e ? Cu E? .34 V standard
  reduction potential
• Cu ? Cu2 2 e E? -.34 V standard
  oxidation potential
• Zn2 2 e ? Zn E? -.76 V standard
  reduction potential
• Zn ? Zn2 2 e E? .76 V standard
  oxidation potential
• Cu2 Zn ? Cu Zn2 Eo .34 V .76
  V 1.1 V

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• Standard electrode potentials are listed in
  tables
• The half-cell reaction are written as reductions
  rather than as
• oxidation
• The oxidation half-cell potential is the negative
  of the reduction
• half-cell potential
• The half-reactions are listed in order of
  decreasing standard
• reduction potential which means
• decreasing tendency to occur in the forward
  direction
• increasing tendency to occur in the reverse
  direction
• Strongest oxidizing agent are in the upper left
  of the table
• strongest reducing agents are in the lower right
  of the table

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Standard Reduction Potential Table

• 1. The half-reactions are all written as
  reductions. Oxidizing agents and electrons are
  on the left side of each half-reaction and
  reducing agents are on the right side
• 2. The listed half-cell potentials are standard
  reduction potentials, also known as standard
  electrode potentials
• 3. The half-reactions are listed in order of
  decreasing standard reduction potential(decreasin
  g tendency to occur in the forward direction
  increasing tendency to occur in the reverse
  direction) The strongest oxidizing agents are
  located in the upper left of the table (F2,
  H2O2, MNO4-) the strongest reducing agents are
  found in the lower right of the table (Li, Na,
  Mg

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Using Standard Reduction Potentials

• Table arranges oxidizing or reducing agents in
  order of increasing strength, this allows us to
  predict the spontaneity or non-spontaneity of
  thousands of redox reactions.
• Lets calculate E? for the oxidation of Zn(s) by
  Ag(aq)
• 2 Ag(aq) Zn(s) ? 2 Ag(s) Zn2(aq)

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Using Standard Reduction Potentials

• Step 1 Find half-reactions in Table and write
  them in the appropriate direction
• Step 2 Multiply the Ag/Ag half-reaction by a
  factor of 2 so that the electrons cancel. Do
  not multiply the E? by 2 because electric
  potential is an intensive property, which does
  not depend on the amount of substance.
• ?E? -?G? / nF
• ?G will double and n will double so that ?E?
  will remain constant
• Step 3 Change E? for oxidations since the table
  is based on reductions
• Step 4 Add the half-reactions to get the
  overall reaction.

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Using Standard Reduction Potentials

• Reduction 2 x Ag e- ? Ag E? 0.80 V
• Oxidation Zn ? Zn2 2e- E?
  -(-0.76 V)
• Overall reaction
• 2 Ag(aq) Zn(s) ? 2 Ag(s) Zn2(aq)
  E? 1.56 V
• ?G?   -?E?(nF) 2 x 96,500 x 1.56 -3.01 x
  105 J
• Because E? is positive and ?G? is negative,
  oxidation of zinc by Ag is a spontaneous
  reaction under SS conditions. Note Ag can
  oxidize any reducing agent that lies below it in
  the table. Cant oxidize a reducing agent that
  appears above it on the table. Because E? will
  have a negative value.

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Example 3

• Arrange the following oxidizing agents in order
  of increasing strength under SS conditions
• Br2(l) , Fe3(aq) , Cr2O72-(aq)

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Example 4

• Arrange the following reducing agents in order of
  increasing strength under SS conditions
• Al(s) , Na(s) , Zn(s)

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Example 5

• Predict from the Reduction table whether each of
  the following reactions can occur spontaneously
  under SS conditions
• a)      2 Fe3(aq) 2 I-(aq) ? 2 Fe2(aq)
  I2(s)
• b)      3 Ni(s) 2 Al3(aq) ? 3 Ni2(aq)
  2 Al(s)

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Cell potentials under non-standard conditions
The Nernst Equation

• Cell pot. and Gibbs free energy depend on
• temp. and composition of reaction mixture
• concentration of solutes
• partial pressures of gases
• ?G? Standard free-energy
•  
• ?G ?G? RT ln Q Q Reaction quotient
• nFE -nFE? RT ln Q dividing by nF we get
• RT
• E E0 - ----- ln Q
•   nF

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The Nernst Equation

• Nernst Equation E E? - RT ln Q
• nF
• or E E? - 2.303 RT log Q
• nF
•  
• or E E? - 0.0592 V log Q in Volts _at_ 25?C
• n
• This allows us to calculate cell potentials under
  non-standard state conditions

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• Consider a galvanic cell that uses the reaction
•  
• Zn(s) 2 H(aq) ? Zn2(aq) H2(g)
•  
• Calculate the cell potential _at_ 25?C when H
  1.0 M, Zn2 0.0010 M, and PH2 0.10 atm
• First step calculate standard cell potential E0
  E?ox E?red
• Second step use Nernst equation to find cell
  potential under the cited conditions
• E E? - 0.0592 V log Q
• n

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First step Eo Eo(Zn EoH2 - (-.076 V)
0.00 V 0.76 V Second step 0.059
V 0.059 Zn2 pH2 E Eo - ----------
log Q 0.76 V - -------- log-------------
n n H E 0.76 V 0.12
V 0.88 V
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Example 7

• Consider a galvanic cell that uses the reaction
• Cu(s) 2 Fe3(aq) ? Cu2(aq) 2 Fe2(aq)
•  
• What is the potential of a cell at 25?C that has
  the following ion concentrations
•  
• Fe3 1.0 x10-4 M Cu2 0.25 M
  Fe2 0.20 M
•  
• Going to use E? E?ox E?red
•  
• Going to use E E? - 0.0592 V log Q
• n

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Eo EoCu EoFe2 - 0.34 V 0.77 V
0.43 V 0.0592 Cu2Fe22 E Eo -
-------- log --------------
n Fe32 0.25 x 0.22 E 0.43 V -
0.0296 log ---------------- (1 x
10-4)2 E 0.43 V - 0.18 V 0.25 V
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Electrochemical determination of pH with a pH
meter voltmeter We need a pH sensitive
electrode glass electrode H2 2
H 2 e E0 0.00 V Nernst
equation E E0 - 0.0592 log H2
------------- --------
n 2, log H2 2 logH n
pH2 pH2 1 atm E 0 V
0.0592 pH
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H2 has a standard electrode potential of 0.00
V which means under standard state
condition H 1 mol/L pH 0 All metals with
a standard potential gt 0.00 V will dissolve in
strong acids (Pb, Ni, Cd, Fe, Zn, Al etc Cu
will not dissolve) In pure water (pH 7) the
standard oxidation potential for H2 is E E0
0.0592 pH 0.00 V 0.414 V 0.414 V All
metals with a standard potential gt 0.414 V will
dissolve in water (Fe, Zn, Al, Mg etc Ni, Cd,
Pb will not dissolve) Why does Al and Mg not
dissolve in water?
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Standard Cell Potentials and Equilibrium
Constants DG0 -n F E0 and DG0 -RT ln K
K equilibrium constant - n F E0 -RT
ln K RT 0.0592 E0 ------- ln K
---------- log K in Volts _at_ 250C n F
n Most common use of this equation is the
calculation of K Zn Cu2 Zn2
Cu E0 1.10 V log K n
E0/0.0592 2 x 1.10 /0.059 37.1 K antilog
37.2 2 x 1037
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Standard Cell potentials and equilibrium constants

• Now we have 3 ways to find an equilibrium
  constant
• 1.    aA bB ? cC dD from concentration or
  partial pressures
• K Cc Dd
• Aa Bb
• 2. ?G? -RT ln K, ln K -?G?
• 
  RT
• From thermochemical data
•  
• 3. E? RT ln K or ln K nFE? or
  log K nE?
• nF
  RT .0592 V
•  
• From electrochemical data

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Example 8

• Use the standard reduction potentials in Table to
  calculate the equilibrium constant at 25?C for
  the reaction
• 6 Br-(aq) Cr2O72-(aq) 14 H(aq) ? 3
  Br2(l) 2 Cr3(aq) 7 H2O(l)

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Example 9

• Use the data in the standard reduction table to
  calculate the equilibrium constant _at_ 25?C for the
  reaction
•  
• 4 Fe2(aq) O2(g) 4 H(aq) ? 4 Fe3(aq)
  2 H2O(l)

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Batteries

• Batteries are the most important practical
  application of galvanic cells.
• Single-cell batteries consist of one galvanic
  cell.
• Multicell batteries consist of several galvanic
  cells linked in series to obtain the desired
  voltage. (car batteries)

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Batteries

• Lead Storage Battery A typical 12 volt battery
  consists of six individual cells connected in
  series.
• Anode Lead grid packed with spongy lead.
• Cathode Lead grid packed with lead (IV) oxide.
• Electrolyte 38 by mass sulfuric acid.
• Cell Potential 1.924 V

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Anode Pb HSO4-
PbSO4 H 2e Cathode
PbO2 3H HSO4- 2e PbSO4
2H2O --------------------------------------------
------------------------------------- Sum Pb
PbO2 2H 2HSO4- 2PbSO4
2H2O Both electrodes (Pb and PbO2) are
solids The reaction product PbSO4 is formed on
both electrodes PbSO4 is insoluble in aqueous
solution and here in sulfuric acid PbSO4 adheres
to the surface of each electrode Recharging the
battery is the opposite process applying a
voltage greater than 2Volt/cell.
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Batteries

• Zinc Dry-Cell Also called a Leclanché cell, uses
  a viscous paste rather than a liquid solution.
• Anode Zinc metal can on outside of cell.
• Cathode MnO2 and carbon black paste on graphite.
• Electrolyte NH4Cl and ZnCl2 paste in starch.
• Cell Potential 1.5 V but deteriorates to 0.8 V
  with use.

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Dry-cell reactions Anode
Zn Zn2
2e Cathode 2MnO2 2NH4 2e
Mn2O3 2NH3 H2O The dry-cell reaction
cannot be recharged Disadvantage of dry-cells
Zn corrodes under acidic conditions, under basic
conditions Zn corrodes much less
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• Alkaline Dry-Cell Modified Leclanché cell which
  replaces NH4Cl with NaOH or KOH.
• Anode Zinc metal can on outside of cell.
• Cathode MnO2 and carbon black paste on graphite.
• Electrolyte NaOH or KOH, and Zn(OH)2 paste.
• Cell Potential 1.5 V but longer lasting, higher
  power, and more stable current and voltage.

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Batteries

• Mercury Dry-Cell Modified Leclanché cell which
  replaces MnO2 with HgO and uses a steel cathode.
• Anode Zinc metal can on outside of cell.
• Cathode HgO in contact with steel.
• Electrolyte KOH, and Zn(OH)2 paste.
• Cell Potential 1.3 V with small size, longer
  lasting, and stable current and voltage.

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Mercury battery Anode Zn
2OH- ZnO 2e Cathode
HgO H2O 2e Hg
2OH- Not rechargeble Recycling problem!
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Batteries

• NickelCadmium Battery Modified Leclanché cell
  which is rechargeable.
• Anode Cadmium metal.
• Cd(s) 2 OH(aq) ? Cd(OH)2(s) 2 e
• Cathode Nickel(III) compound on nickel metal.
• NiO(OH) (s) H2O(l) e ? Ni(OH)2(s) OH(aq)
• Electrolyte Nickel oxyhydroxide, NiO(OH).
• Cell Potential 1.30 V

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Batteries

• NickelMetalHydride (NiMH)
• Replaces toxic Cd anode witha hydrogen atom
  impregnated ZrNi2 metal alloy.
• During oxidation at the anode, hydrogen atoms are
  released as H2O.
• Recharging reverses this reaction.

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Batteries

• Lithium Ion (Liion) The newest rechargeable
  battery is based on the migration of Li ions.
• Anode Li metal, or Li atom impregnated graphite.
• Li(s) ? Li e
• Cathode Metal oxide or sulfide that can accept
  Li.
• MnO2(s) Li(aq) e ? LiMnO2(s)
• Electrolyte Lithium-containing salt such as
  LiClO4, in organic solvent. Solid state polymers
  can also be used.
• Cell Potential 3.0 V

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• Fuel Cell
• H2 reacts with O2 to form water.
• The two half-cell reactions, the oxidation of H2
  and the reduction
• of O2, take place in two different compartments
  separated with a
• membrane.
• Anode Porous carbon containing metallic
  catalysts.
• 2 H2(s) 4 OH(aq) ? 4 H2O(l) 4 e
• Cathode Porous carbon containing metallic
  catalysts.
• O2(s) 2 H2O(l) 4 e ? 4 OH(aq)
• Electrolyte Hot aqueous KOH solution.
• Cell Potential 1.23 V, but is only 40 of cell
  capacity.

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Fuel Cell

• Fuel cells are not batteries because they are not
  self-contained.
• Fuel cells typically have about 40 conversion to
  electricity the remainder is lost as heat.
• Excess heat can be used to drive turbine
  generators.

56
Corrosion

• Corrosion is the oxidative deterioration of a
  metal, such as iron to rust.
• To prevent corrosion first we have to understand
  how it occurs
• a) Rusting requires both oxygen and water
• b) Rusting results from tiny galvanic cells
  formed by water droplets

57
Rust has approximately composition
Fe2O3H2O Rusting is an electrochemical reaction
Two different regions within a water drop, an
anodic and a cathodic region
58
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) ? 2
Fe2(aq) 2 H2O(l)

• This electrochemical mechanism for corrosion also
  explains why automobiles rust more rapidly in
  parts of the country where road salt is used to
  melt snow and ice. Dissolved salts in the water
  droplet greatly increase the conductivity of the
  electrolyte, thus accelerating the pace of
  corrosion.

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Prevention of corrosion

• Galvanizing is the coating of iron with zinc.
  Zinc is more easily oxidized than iron, which
  protects and reverses oxidation of the iron.
• Cathodic Protection is the protection of a metal
  from corrosion by connecting it to a metal (a
  sacrificial anode) that is more easily oxidized.
• All that is required is an electrical connection
  to the sacrificial anode (usually magnesium or
  zinc).

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Prevention of corrosion

• Cathodic Protection with Zinc Layer

61
Prevention of corrosion

• Cathodic Protection with Magnesium Anode

62
Electrolysis

• Electrolysis is the process in which electrical
  energy is used to drive a nonspontaneous chemical
  reaction.
• An electrolytic cell is an apparatus for carrying
  out electrolysis.
• Processes in an electrolytic cell are the reverse
  of those in a galvanic cell.

63
Applications of Electrolysis

• Manufacture of Sodium (Downs Cell)

64
Applications of Electrolysis

• Manufacture of Cl2 and NaOH (ChlorAlkali)

65
Applications of Electrolysis

• Manufacture of Aluminum (HallHeroult)

66
Applications of Electrolysis

• Electrorefining and Electroplating