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Chapter 19
Electrochemistry
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19.1 Balancing Redox Reactions
• Half-reaction method – acid solution
– Can add H2O
– Can add H+
• Steps
– Separate the unbalanced reaction into halfreactions.
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– Balance each of the half-reactions with regard
to atoms other than O and H.
– Balance both half-reactions for O by adding
H2O.
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– Balance both half-reactions for H by adding
H+.
– Balance both half-reactions for charge by
adding electrons.
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– multiply one or both of the half-reactions
by the number(s) required to make the number
of electrons the same in both.
– add the balanced half-reactions together and
cancel any identical terms that appear on both
sides.
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• Additional steps needed in basic solution
– For each H+ ion in the final equation, add one
OH ion to each side of the equation,
combining the H+ and OH ions to produce
H2O.
– Make any additional cancellations made
necessary by the new H2O molecules.
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19.2 Galvanic Cells
• Galvanic cell - the experimental
apparatus for generating electricity
through the use of a spontaneous reaction
• Electrodes
– Anode (oxidation)
– Cathode (reduction)
• Half-cell - combination of container,
electrode and solution
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• Salt bridge - conducting medium through
which the cations and anions can move
from one half-cell to the other.
• Ion migration
– Cations – migrate toward the cathode
– Anions – migrate toward the anode
• Cell potential (Ecell) – difference in
electrical potential between the anode and
cathode
– Concentration dependent
– Temperature dependent
– Determined by nature of reactants
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• cell diagram convention
anode
cathode
salt bridge
phase boundary
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19.3 Standard Reduction Potentials
• Designated Eo
• Measured relative to the standard
hydrogen electrode (SHE)
– Standard conditions
– Assigned a value of 0 V
– Reaction
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o
o
o
Ecell
 Ecathode
 Eanode
o
o
Ecell
 EHo /H  E Zn
2
/Zn
2
o
0.76V  0 V  EZn
2
/Zn
o
EZn
 0.76V
2
/Zn
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o
o
o
Ecell
 Ecathode
 Eanode
o
o
o
Ecell
 ECu

E
2
/Cu
H /H
2
o
0.34 V  ECu
0 V
2
/Cu
o
ECu
 0.34 V
2
/Cu
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o
o
o
Ecell
 Ecathode
 Eanode
o
o
o
Ecell
 ECu

E
 0.34V  (0.76V)
2
/Cu
Zn2 /Zn
o
Ecell
 1.10 V
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Determine the overall cell reaction and
E°cell (at 25°C) of a galvanic cell made of
an Al electrode in a 1.0 M Al(NO3)3 solution
and a Cu electrode in a 1.0 M Cu(NO3)2
solution.
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Cell Potential
Anode
Cathode
3+
Al
+
2+
Cu +
-
3e
-
2e
o
1.66 V
Al
E=
Cu
Eo= + 0.34 V
o
o
o
Ecell
 Ecathode
 Eanode
o
o
o
Ecell
 ECu

E
2
/Cu
Al3 /Al
o
Ecell
 0.34 V  ( 1.66 V)  2.00 V
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Cell Reaction
3Cu2+ + 6e2Al
2Al
+
3Cu
3+
2Al
2+
3Cu
3+
2Al
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6e
+ 3Cu
+
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19.4 Spontaneity of Redox Reactions
Under Standard-State Conditions
• Based on electric charge and work
G  nFEcell
– n is the number of moles of electrons
– F is the Faraday constant
• 96,500 J/V . mol e
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• At standard conditions
o
G o  nFEcell
• Relation to the equilibrium constant
RT
o
Ecell 
ln K
nF
– Using the values of R, T (298 K) and F
E
o
cell
0.0257 V

ln K
n
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Determine the value of a) Go and b) K for
the following reaction.
2Al
+
2+
3Mn
3+
2Al
+
3Mn
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o
Ecell
 0.48 V
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a) Standard free energy change, Go
G  nFE
o
o
cell


G  (6 mol e )(96,500 J/V  mol e )(0.48 V)
o
G  2.78 x10 J/mol
o
5
G o  2.78x10 2 kJ/mol
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b) Equilibrium constant, K
E
o
cell
0.0257 V

ln K
n
K e
K e
nE o
0.0257V
(6)(0.48V)
0.0257V
 4.7x10 48
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19.5 Spontaneity of Redox Reactions
Under Conditions Other than StandardState Conditions
• The Nernst Equation –derived from
thermodynamics
0.0592 V
E E 
log Q
n
o
• or
0.0257 V
E E 
ln Q
n
o
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Will the following reaction occur spontaneously at
298 K if [Fe2+] = 0.60 M and [Cd2+] = 0.010 M?
Cd(s)
+
2+
Fe (aq)
Cd2+(aq)
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+
Fe(s)
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Cd(s)
+
2+
Fe (aq)
Cd2+(aq)
+
Fe(s)
o
o
o
Ecell
 Ecathode
 Eanode
o
o
o
Ecell
 EFe/Fe

E
2
Cd/Cd2
o
Ecell
 0.44 V  ( 0.40 V)  0.04 V
0.0257 V
E E 
ln Q
n
o
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2
0.0257 V [Cd ]
E E 
ln
2
n
[Fe ]
o
0.0257 V (0.010 M )
E  (0.04 V) 
ln
2
(0.60 M )
E  0.01 V
Spontaneous since E is a positive even though very
small.
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• Concentration cells – galvanic cells composed of the
same material but differing in ion concentrations
0.0257 V
[Dilute ]
E E 
ln
n
[Concentrat ed]
o
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19.6 Batteries
• A battery is a galvanic cell, or a series of
cells connected that can be used to deliver
a self-contained source of direct electric
current.
• Dry Cells and Alkaline Batteries
– no fluid components
– Zn container in contact with MnO2 and an
electrotyte
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Dry Cell
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Alkaline Cell
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• Lead Storage Batteries
– Six identical cells in series
– Lead anode and PbO2 cathode
– Immersed in H2SO4
– Each cell delivers ~ 2 V
– Rechargeable
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Lead Storage Battery
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• Lithium-Ion Batteries
– The overall cell potential is 3.4 V, which is a
relatively large potential.
– Lithium is also the lightest metal—only
6.941 g of Li (its molar mass) are needed to
produce 1 mole of electrons.
– recharged hundreds of times.
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• Fuel Cells
– Direct production of electricity by
electrochemical means
– Increased efficiency of power production
– In its simplest form
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Simple Form of Fuel Cell
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19.7 Electrolysis
• electrolysis - the use of electric energy to
drive a nonspontaneous chemical reaction
• electrolytic cell – the cell used to carry
out electrolysis
– same principles apply to both galvanic and
electrolytic cells
– in aqueous solutions you must also consider
the oxidation or reduction of water
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• Examples
– Molten sodium chloride
– Carried out in a Downs cell
Laboratory
Commercial
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– Electrolysis of water
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– Electrolysis of aqueous sodium chloride solution
• Possible anode reactions
• Possible cathode reactions
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• Consider overvoltage – difference between
estimated and actual voltage
–Overvoltage for O2 is higher than Cl2
• Cell reaction
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• Quantitative Applications- measure the
current (in amperes) that passes through
an electrolytic cell in a given period of
time.
– Mass of product formed or reactant
consumed
• Proportional to the amount of
electricity transferred
• Molar mass
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–Defining equation
Charge  Current x time
C  A s
Coulomb (C)
Second (s)
Ampere (A)
(coulomb/second)
C
96,500
mol e 
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A constant current of 0.912 A is passed
through an electrolytic cell containing molten
MgCl2 for 18 h. What mass of Mg is produced?
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Mg2++ 2e-
Mg
 3600 s  1 C 
4
coulombs  (0.912 A)18 h 

  5.91 x 10 C
 h  A  s 


 1mol Mg 
1
mol
e
4

mol Mg  (5.91 x 10 C)
 0.306 mol Mg
 
 96,500 C  2 mol e 
 24.31 g Mg 
  7.44 g Mg
g Mg  0.306 mol Mg
 mol Mg 
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19.8 Corrosion
• Corrosion - generally refers to the
deterioration of a metal by an
electrochemical process.
• Many metals undergo corrosion
– For example, corrosion of Fe, oxidation of Al
• Can be enhanced by atmospheric
conditions ( e.g. acidic medium)
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Electrochemical Process for the Formation of Rust
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• Electochemical processes can be used to
prevent corrosion
– Passivation – formation of a thin oxide layer
by treating with an oxidizing agent
– Formation of an alloy
– Coating with a layer of a less active metal
• Tin cans
– Galvanization (zinc-plating)
• Zinc oxide coating constitutes the
protective coating
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Key Points
• Balancing redox reactions
– Half-reaction method
• Galvanic cells
– Electrodes (anode and cathode)
– Cell potential (Ecell)
– Cell diagrams
• Standard Reduction potentials
– Relative to the standard hydrogen electrode
– Used to determine cell potential
o
o
o
Ecell
 Ecathode
 Eanode
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• Spontaneity – standard conditions
– Free energy
o
G o  nFEcell
– Equilibrium Constant
E
o
cell
0.0257 V

ln K
n
• Spontaneity – nonstandard conditions
– Nernst Equation
0.0592 V
E E 
log Q
n
o
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• Batteries
– Dry cell and alkaline batteries
– Lead storage battery
– Lithium-ion batteries
– Fuel cells
• Electrolysis
– Molten salts
– Aqueous solutions
– Quantitative applications
• Corrosion
– Metal deterioration
– Prevention
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