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7/3/2012
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Electrochemistry
© 2009, Prentice-Hall, Inc.
Chapter 20
Electrochemistry
Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
John D. Bookstaver
St. Charles Community College
Cottleville, MO
Electrochemistry
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Chapter 20 Problems
• Problems 15, 17, 19, 23, 27, 29, 33, 39,
59
Electrochemistry
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Electrochemistry
In the broadest sense, electrochemistry
is the study of chemical reactions that
produce electrical effects and of the
chemical phenomena that are caused
by the action of currents or voltages.
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Electrochemistry
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Oxidation-Reduction Reactions
Basic Terms
Oxidation - loss of electrons
Reduction - gain of electrons
Redox reaction
oxidizing agent - substance that causes
oxidation by being reduced
reducing agent - substance that causes
reduction by being oxidized
Electrochemistry
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Electrochemical Reactions
In electrochemical reactions, electrons
are transferred from one species to
another.
Electrochemistry
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Oxidation Numbers (State)
In order to keep
track of what loses
electrons and what
gains them, we
assign oxidation
numbers.
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Electrochemistry
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Oxidation and Reduction
• A species is oxidized when it loses electrons.
– Here, zinc loses two electrons to go from neutral
zinc metal to the Zn2+ ion.
Electrochemistry
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Oxidation and Reduction
• A species is reduced when it gains electrons.
– Here, each of the H+ gains an electron, and they
combine to form H2.
Electrochemistry
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Oxidation and Reduction
• What is reduced is the oxidizing agent.
– H+ oxidizes Zn by taking electrons from it.
• What is oxidized is the reducing agent.
– Zn reduces H+ by giving it electrons.
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Electrochemistry
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Oxidation and Reduction
• Oxidation
– O.S. of some element increases in the
reaction.
– Electrons are on the right of the equation
• Reduction
– O.S. of some element decreases in the
reaction.
– Electrons are on the left of the equation.
Electrochemistry
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Assigning Oxidation Numbers
1. Elements in their elemental form have
an oxidation number of 0.
2. The oxidation number of a monatomic
ion is the same as its charge.
Electrochemistry
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Assigning Oxidation Numbers
3. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
– Oxygen has an oxidation number of −2, except in the peroxide ion, which has an oxidation number of −1.
– Hydrogen is −1 when bonded to a metal and +1 when bonded to a nonmetal.
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Electrochemistry
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Assigning Oxidation Numbers
3. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
– Fluorine always has an oxidation number of −1.
– The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.
Electrochemistry
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Assigning Oxidation Numbers
4. The sum of the oxidation numbers in a
neutral compound is 0.
5. The sum of the oxidation numbers in a
polyatomic ion is the charge on the
ion.
Electrochemistry
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Oxidation State
What is the oxidation state of S in H2SO4?
• H => +1
• O => -2
• neutral compound, thus sum equals
zero
• 4O => 4*-2 = -8
• 2H => 2*+1 = +2
• 0 = +2 + (x) + (-8)
x = +6
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Electrochemistry
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Oxidation State
What is the oxidation state of Cl in
HClO4?
• H => +1
• O => -2
• neutral compound, thus sum equals
zero
• 4O => 4*-2 = -8
• H => 1*+1 = +1
• 0 = +1 + (y) + (-8)
y = +7
Electrochemistry
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Oxidation and Reduction Half-
Reactions • A reaction represented by two half-
reactions.
Oxidation:
Reduction:
Overall:
Zn(s) → Zn2+(aq) + 2 e-
Cu2+(aq) + 2 e- → Cu(s)
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Electrochemistry
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Balancing Oxidation-Reduction
Equations
• Few can be balanced by inspection.
• Systematic approach required.
• The Half-Reaction (Ion-Electron)
Method
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Electrochemistry
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Balancing Oxidation-Reduction
Equations
Perhaps the easiest way to balance the
equation of an oxidation-reduction
reaction is via the half-reaction method.
Electrochemistry
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Balancing Oxidation-Reduction
Equations
This involves treating (on paper only) the
oxidation and reduction as two separate
processes, balancing these half reactions,
and then combining them to attain the
balanced equation for the overall reaction.
Electrochemistry
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The Half-Reaction Method
1. Assign oxidation numbers to
determine what is oxidized and what is
reduced.
2. Write the oxidation and reduction half-
reactions.
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Electrochemistry
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The Half-Reaction Method
3. Balance each half-reaction.
a. Balance elements other than H and O.
b. Balance O by adding H2O.
c. Balance H by adding H+.
d. Balance charge by adding electrons.
4. Multiply the half-reactions by integers so that the electrons gained and lost are the same.
Electrochemistry
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The Half-Reaction Method
5. Add the half-reactions, subtracting things that appear on both sides.
6. Make sure the equation is balanced according to mass.
7. Make sure the equation is balanced according to charge.
Electrochemistry
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The Half-Reaction Method
Consider the reaction between MnO4− and C2O4
2− :
MnO4−
(aq) + C2O42−
(aq) Mn2+ (aq) + CO2 (aq)
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Electrochemistry
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The Half-Reaction Method
First, we assign oxidation numbers.
MnO4− + C2O4
2- Mn2+ + CO2
+7 +3 +4 +2
Since the manganese goes from +7 to +2, it is reduced.
Since the carbon goes from +3 to +4, it is oxidized.
Electrochemistry
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Oxidation Half-Reaction
C2O42− CO2
To balance the carbon, we add a
coefficient of 2:
C2O42− 2 CO2
Electrochemistry
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Oxidation Half-Reaction
C2O42− 2 CO2
The oxygen is now balanced as well. To
balance the charge, we must add 2
electrons to the right side.
C2O42− 2 CO2 + 2 e−
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Electrochemistry
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Reduction Half-Reaction
MnO4− Mn2+
The manganese is balanced; to balance
the oxygen, we must add 4 waters to
the right side.
MnO4− Mn2+ + 4 H2O
Electrochemistry
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Reduction Half-Reaction
MnO4− Mn2+ + 4 H2O
To balance the hydrogen, we add 8 H+
to the left side.
8 H+ + MnO4− Mn2+ + 4 H2O
Electrochemistry
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Reduction Half-Reaction
8 H+ + MnO4− Mn2+ + 4 H2O
To balance the charge, we add 5 e− to
the left side.
5 e− + 8 H+ + MnO4− Mn2+ + 4 H2O
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Electrochemistry
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Combining the Half-Reactions
Now we evaluate the two half-reactions
together:
C2O42− 2 CO2 + 2 e−
5 e− + 8 H+ + MnO4− Mn2+ + 4 H2O
To attain the same number of electrons
on each side, we will multiply the first
reaction by 5 and the second by 2.
Electrochemistry
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Combining the Half-Reactions
5 C2O42− 10 CO2 + 10 e−
10 e− + 16 H+ + 2 MnO4− 2 Mn2+ + 8 H2O
When we add these together, we get:
10 e− + 16 H+ + 2 MnO4− + 5 C2O4
2−
2 Mn2+ + 8 H2O + 10 CO2 +10 e−
Electrochemistry
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Combining the Half-Reactions
10 e− + 16 H+ + 2 MnO4− + 5 C2O4
2−
2 Mn2+ + 8 H2O + 10 CO2 +10 e−
The only thing that appears on both sides are the electrons. Subtracting them, we are left with:
16 H+ + 2 MnO4− + 5 C2O4
2−
2 Mn2+ + 8 H2O + 10 CO2
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Electrochemistry
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Balancing in Basic Solution
• If a reaction occurs in basic solution, one
can balance it as if it occurred in acid.
• Once the equation is balanced, add OH−
to each side to “neutralize” the H+ in the
equation and create water in its place.
• If this produces water on both sides, you
might have to subtract water from each
side.
Electrochemistry
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Important Electrochemical Terms
An electrochemical cell is a device that combines
two half-cells with the appropriate connections
between electrodes and solutions
A voltaic cell is an electrochemical cell in which
electric current is generated from a spontaneous
redox reaction
The anode is the electrode at which oxidation occurs
and the cathode is the electrode at which reduction
occurs
EOS
The cell potential (Ecell) is the potential difference
that propels electrons from the anode to the cathode
Electrochemistry
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Voltaic Cells
In spontaneous
oxidation-reduction
(redox) reactions,
electrons are
transferred and
energy is released.
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Electrochemistry
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Cell Diagrams – Conventions
Place the anode on the left side of the diagram
Place the cathode on the right side of the diagram
Use a single vertical line ( | ) to represent the
boundary between different phases, such as
between an electrode and a solution
EOS
Use a double vertical line ( || ) to represent a salt
bridge or porous barrier separating two half-
cells
Electrochemistry
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An Example Cell Diagram
EOS
Electrochemistry
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Terminology
• Galvanic (Voltaic) cells.
– Produce electricity as a result of
spontaneous reactions.
• Electrolytic cells.
– Non-spontaneous chemical change driven
by electricity.
• Couple, M|Mn+
– A pair of species related by a change in
number of e-.
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Electrochemistry
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Voltaic Cells
• harnessed chemical reaction which
produces an electric current
Electrochemistry
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Terminology
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Ecell = 1.103 V
Electrochemistry
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Voltaic Cells
• We can use that
energy to do work if
we make the
electrons flow
through an external
device.
• We call such a
setup a voltaic cell.
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Electrochemistry
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Voltaic Cells
• A typical cell looks
like this.
• The oxidation
occurs at the anode.
• The reduction
occurs at the
cathode.
Electrochemistry
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Voltaic Cells
Once even one
electron flows from
the anode to the
cathode, the
charges in each
beaker would not be
balanced and the
flow of electrons
would stop.
Electrochemistry
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Voltaic Cells
• Therefore, we use a
salt bridge, usually a
U-shaped tube that
contains a salt
solution, to keep the
charges balanced.
– Cations move toward
the cathode.
– Anions move toward
the anode.
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Electrochemistry
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Voltaic Cells • In the cell, then,
electrons leave the
anode and flow
through the wire to
the cathode.
• As the electrons
leave the anode,
the cations formed
dissolve into the
solution in the
anode
compartment.
Electrochemistry
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Voltaic Cells • As the electrons
reach the cathode,
cations in the
cathode are
attracted to the now
negative cathode.
• The electrons are
taken by the cation,
and the neutral
metal is deposited
on the cathode.
Electrochemistry
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Electromotive Force (emf)
• Water only
spontaneously flows
one way in a
waterfall.
• Likewise, electrons
only spontaneously
flow one way in a
redox reaction—from
higher to lower
potential energy.
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Electrochemistry
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Electromotive Force (emf)
• The potential difference between the
anode and cathode in a cell is called the
electromotive force (emf).
• It is also called the cell potential and is
designated Ecell.
Electrochemistry
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Cell Potential
Cell potential is measured in volts (V).
1 V = 1 J
C
Electrochemistry
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Cell Potential (Ecell)
The cell potential (Ecell) is the potential
difference that propels electrons from
the anode to the cathode
E°cell = Eocathode - E
oanode
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Electrochemistry
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Standard Electrode Potentials
• Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements.
• The potential of an individual electrode is difficult to establish.
• Arbitrary zero is chosen for the Standard Hydrogen Electrode (SHE)
Electrochemistry
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Standard Hydrogen Electrode
• Half-cell values are referenced to a standard
hydrogen electrode (SHE).
• By definition, the reduction potential for
hydrogen is 0 V:
2 H+ (aq, 1M) + 2 e− H2 (g, 1 atm)
Electrochemistry
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Standard Hydrogen Electrode
In the standard
hydrogen electrode,
hydrogen gas at exactly
1 bar pressure is
bubbled over an inert
platinum electrode and
into an aqueous solution
with the concentration
adjusted so that the
activity of H3O+ is
exactly equal to one EOS
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Electrochemistry
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Standard Hydrogen Electrode 2 H+(a = 1) + 2 e- H2(g, 1 bar) E° = 0
V
Pt|H2(g, 1 bar)|H+(a = 1)
Electrochemistry
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Standard Electrode Potential, E°
• E° defined by international agreement.
• The tendency for a reduction process to
occur at an electrode.
– All ionic species present at a=1
(approximately 1 M).
– All gases are at 1 bar (approximately 1
atm).
– Where no metallic substance is indicated,
the potential is established on an inert
metallic electrode (ex. Pt).
Electrochemistry
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Standard Electrode Potentials
A standard electrode potential, Eo, is based on
the tendency for reduction to occur at the electrode
The cell voltage, called the standard cell potential
(Eocell), is the difference between the standard
potential of the cathode and that of the anode
EOS
Two illustrations … Cu and Zn
Voltaic cells can produce electrical work
w = –n F Ecell
VideoClip
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Electrochemistry
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Standard Cell Potential Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V
E°cell = E°cathode - E°anode
E°cell = E°Cu2+/Cu - E°H+/H2
0.340 V = E°Cu2+/Cu - 0 V
E°Cu2+/Cu = +0.340 V
H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s) E°cell = 0.340
V
Electrochemistry
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Measuring Standard Reduction
Potential
anode anode cathode cathode
Electrochemistry
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Cu2+/Cu Electrode
EOS
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Electrochemistry
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Zn2+/Zn Electrode
EOS
Electrochemistry
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Observed Voltages …
EOS
Electrochemistry
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Standard Reduction Potentials
Reduction
potentials for
many
electrodes
have been
measured and
tabulated.
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Electrochemistry
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Important Points About
Electrode and Cell Potentials
Electrode potentials and cell voltages are intensive
properties – independent of the amount of matter
Cell voltages can be ascribed to oxidation–
reduction reactions without regard to voltaic cells
EOS
Ecell = Ecathode – Eanode
Electrochemistry
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Cell Potentials
• For the oxidation in this cell,
• For the reduction,
Ered = −0.76 V
Ered = +0.34 V
Electrochemistry
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Cell Potentials
Ecell = Ered (cathode) − Ered (anode)
= +0.34 V − (−0.76 V)
= +1.10 V
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Electrochemistry
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Oxidizing and Reducing Agents
• The strongest
oxidizers have the
most positive
reduction potentials.
• The strongest
reducers have the
most negative
reduction potentials.
Electrochemistry
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Oxidizing and Reducing Agents
The greater the
difference between
the two, the greater
the voltage of the
cell.
Electrochemistry
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Criteria for Spontaneous
Change
If Ecell is positive, the reaction in the forward
direction (from left to right) is spontaneous
If Ecell is negative, the reaction is nonspontaneous
If Ecell = 0, the system is at equilibrium
EOS
When a cell reaction is reversed, Ecell and DG
change signs
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Electrochemistry
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Free Energy
DG for a redox reaction can be found by
using the equation
DG = −nFE
where n is the number of moles of
electrons transferred, and F is a
constant, the Faraday.
1 F = 96,485 C/mol = 96,485 J/V-mol
Electrochemistry
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Free Energy
Under standard conditions,
DG = −nFE
Electrochemistry
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Nernst Equation
• Remember that
DG = DG + RT ln Q
• This means
−nFE = −nFE + RT ln Q
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Electrochemistry
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Nernst Equation
Dividing both sides by −nF, we get the
Nernst equation:
E = E − RT
nF ln Q
or, using base-10 logarithms,
E = E − 2.303 RT
nF log Q
Electrochemistry
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Nernst Equation
At room temperature (298 K),
Thus the equation becomes
E = E − 0.0592
n log Q
2.303 RT
F = 0.0592 V
Electrochemistry
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Summary of Important
Relationships
EOS
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Electrochemistry
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Concentration and Cell
Voltage If the discussion is limited to 25 oC, the equation is
o ocell cell cell
0.0592ln ln
RTE E Q E Q
nF n
The Nernst equation relates a cell voltage for
nonstandard conditions (Ecell) to a standard cell voltage,
Eocell, and the concentrations of reactants and products
EOS
The Nernst equation is useful for determining the
concentration of a species in a voltaic cell through
a measurement of Ecell
Electrochemistry
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A Sample Problem
ocell cell
0.0592ln E E Q
n
EOS
Electrochemistry
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Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)
Applying the Nernst Equation for Determining Ecell. What
is the value of Ecell for the voltaic cell pictured below and
diagrammed as follows?
EXAMPLE
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Electrochemistry
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Ecell = Ecell° - log Q n
0.0592 V
Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)
Ecell = Ecell° - log n
0.0592 V [Fe3+]
[Fe2+] [Ag+]
Fe2+(aq) + Ag+(aq) → Fe3+(aq) + Ag (s)
Ecell = 0.029 V – 0.018 V = 0.011 V
EXAMPLE
Electrochemistry
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Concentration Cells
If the cell potential is
determined solely by a
difference in the
concentration of solutes
in equilibrium with
identical electrodes, that
cell is called a
concentration cell
EOS
Electrochemistry
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Concentration Cells
• Notice that the Nernst equation implies that a cell
could be created that has the same substance at
both electrodes.
• For such a cell, would be 0, but Q would not. Ecell
• Therefore, as long as the concentrations
are different, E will not be 0.
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Electrochemistry
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pH Measurement
The Nernst equation can be rewritten in terms of
concentration cells using hydrogen electrodes
EOS
Ecell = (0.0592)(pH)
One useful concentration cell is the hydrogen cell
… it is directly related to pH
Electrochemistry
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Applications of
Oxidation-Reduction
Reactions
Electrochemistry
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Batteries
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Electrochemistry
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Alkaline Batteries
Electrochemistry
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Hydrogen Fuel Cells
Electrochemistry
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Corrosion and…
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Electrochemistry
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…Corrosion Prevention