Electrochemistry
Voltaic Cell (or Galvanic Cell)
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The energy released in a spontaneous redox
reaction can be used to perform electrical
work.
A voltaic cell is a device in which the
transfer of electrons takes place through an
external pathway rather than directly
between reactants.
In a voltaic cell, chemical energy is changed
to electrical energy.
Anode/Cathode
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The two solid metals that are connected by
the external circuit are called electrodes.
The electrode at which oxidation occurs is
called the anode.
The electrode at which reduction occurs is
called the cathode.
Half Cells
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The voltaic cell is thought of as being
comprised of two "half-cells."
One cell is where oxidation occurs and the
other is reduction.
Half Reactions
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Anode (oxidation half-reaction):
Zn (s)
→
Zn+2 (aq) + 2e-

Cathode (reduction half-reaction):
Cu+2 (aq) +2e- → Cu (s)
Determining half reactions will be shown later.
OIL – RIG
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Remember the acronym OIL – RIG
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Oxidation Is Loss of electrons
Reduction Is Gain of electrons
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Red-Ox Reaction
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In a red-ox reaction, one substance must be
oxidized and another substance must be
reduced.
The substance that is oxidized is the
reducing agent.
The substance that is reduced is the
oxidizing agent.
Cell Operation
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Electrons become available when the zinc metal is
oxidized at the anode.
The electrons flow through the external circuit to the
cathode, where they are consumed as Cu+2 is
reduced. Because zinc is oxidized in the cell, the
zinc electrode loses mass, and the concentration of
the Zn+2 solution increases as the cell operates. As
the same time, the Cu electrode gains mass, and the
Cu+2 solution becomes less concentrated as the
Cu+2 is reduced to Cu (s).
Cell Operation
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As the voltaic cell operates, oxidation of Zn
introduces additional Zn+2 ions into the anode
compartment. Unless a means is provided to
neutralize this positive charge, no further
oxidation can take place.
At the same time, the reduction of Cu+2 at the
cathode leaves an excess of negative charge
in solution in that compartment.
Salt Bridge
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Electrical neutrality of the system is
maintained by a migration of ions through the
porous glass disc (salt bridge) that separates
the two compartments.
A salt bridge consists of a U-shaped tube that
contains an electrolyte solution whose ions
will not react with other ions in the cell or with
the electrode materials.
Salt Bridge
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As oxidation and reduction proceed at the
electrodes, ions from the salt bridge migrate
to neutralize charge in the cell
compartments.
Ion Migration
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Anions migrate toward the anode and cations
toward the cathode.
No measurable electron flow will occur
through the external circuit unless a means is
provided for ions to migrate through the
solution from one electrode compartment to
another, completing the circuit.
Electron Flow
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In any voltaic cell the electrons flow from
the anode through the external circuit to
the cathode.
Because the negatively charged electrons
flow from the anode to the cathode, the
anode in a voltaic cell is labeled with a
negative sign and the cathode with a positive
sign.
Oxidation Numbers
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Remind yourself how to determine oxidation
numbers.
Oxidation numbers show what the charge of
each atom would be (in a molecule or ion), if
each atom were an ion.
Go to the text and revisit the Rules for
Assigning Oxidation Numbers section.
Balancing Red-Ox Equations
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Obey the Law of Conservation of Mass
Gain and loss of electrons must be balanced
Half Reactions
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Cu + 2Ag+ →
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Determine the oxidation number for each
substance in the reaction
Write two half reactions based on one
substance being oxidized and one being
reduced.
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2 Ag + Cu+2
Half Reactions
Cu
→
Cu+2
2e- + 2Ag+ →
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+ 2e-
2 Ag
oxidation
reduction
Half reactions show the number of electrons gained
or lost by the substance, standard redox equations
do not.
Go to ChemReview packet to practice half reaction
writing.
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Review Sample Ex: 20.4
Try Practice Exercises
Cell EMF
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What is the driving force that pushes the electrons
through an external circuit in a voltaic cell?
An oxidizing agent in one compartment pulls
electrons through a wire from a reducing agent in the
other compartment.
The pull (or driving force) on the electrons is called
the cell potential ( Ecell ) or electromotive force (emf)
of the cell.
Or, electrons flow from the anode to the cathode (in
a voltaic cell) because of a difference in potential
energy.
Potential Energy of Electrons
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The potential energy of electrons is higher in
the anode than in the cathode. Therefore
electrons spontaneously flow through an
external circuit from the anode to the
cathode.
For any cell reaction that proceeds
spontaneously (i.e. voltaic cell), the cell
potential will be positive.
Standard Conditions
Standard Conditions:
* 1M concentrations for reactants and
products in solution
* 1 atm pressure (for gases)
* 25 ˚C
Under Standard Conditions the emf is called the
standard emf or the standard cell potential (E ˚cell )
Keep in mind that the superscript ˚ denotes standardstate conditions.
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Cell Potential
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The cell potential of a voltaic cell depends on the
particular cathode and anode half-cells involved.
Standard potentials have been assigned to each
individual half-cell, and then use the half-cell
potentials to determine E ˚cell .
The cell potential is the difference between two
electrode potentials (anode and cathode).
The potential associated with each electrode is
chosen to be the potential for reduction to occur at
that electrode.
Standard Reduction Potentials
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Standard electrode potentials are tabulated for reduction reactions, so
they are called standard reduction potentials ( E ˚red ).
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E ˚cell = E ˚red (cathode) - E ˚red (anode)
Or
E ˚cell = E ˚red (cathode) + E ˚ox (anode)
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The hydrogen half-reaction was used as a reference for all of the other
half-reactions.
The standard reduction potential for the hydrogen half-reaction is
assigned a standard reduction potential of exactly 0 V.
It is also called the standard hydrogen electrode. (SHE)
Calculating EMF
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We will use standard reduction potentials to
calculate the emf of a voltaic cell.
Because electrical potential measures
potential energy per electrical charge,
standard reduction potentials are intensive
properties.
Changing the stoichiometric coefficient in a
half-reaction does not affect the value of the
standard reduction potential.
Calculating EMF
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Try sample exercise 20.5
E ˚red and spontaneity
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The more positive the E ˚red value for a halfreaction, the greater the tendency for the
reactant of the half-reaction to be reduced,
and, therefore, to oxidize another species.
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Try sample exercise 20.8
EMF and Free Energy Change
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ΔG = - nFE
n is a positive # that represents the # of etransferred in the reaction
F is Faraday’s constant (this constant is the
quantity of electrical charge on 1 mole of
electrons)
1F = 96,500 C/mol = 96,500 J / V mole
The units for ΔG are J/mole
EMF and Free Energy Change
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ΔG = - nFE
If n and F are positive values, a positive
value of E leads to a negative ΔG.
Remember that a negative ΔG indicates a
spontaneous reaction.
The equation can be altered slightly if the
reactants and products are in their standard
states: ΔG˚ = - nFE˚
Concentration and Cell EMF
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Remember ΔG = ΔG˚ + RT ln Q
Combining
ΔG = ΔG˚ + RT ln Q and ΔG = - nFE
You have the Nernst equation:
E = E˚ – RT ln Q
nF
Nernst Equation
The Nernst equation can be expressed two
ways:
E = E˚ – RT ln Q
nF
E = E˚ – 2.303RT log Q
nF
Nernst Equation Variation
The equation can be simplified if the cell is run
under standard conditions:
E = E˚ – 0.0592 V log Q
n
* Temp. is 298 K
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Use the Nernst Equation to work through the
example on the bottom of page 773
Try the sample exercise 20.11 and practice
exercise
Electrolysis
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This process is the opposite of a voltaic cell.
Ch 20 Problems
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6, 7, 9, 14, 24, 25 a,b, 27, 29, 31, 36, 43, 44,
46, 49, 51, 52, 56, 59