Redox Reactions

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Redox Reactions
Electrochemical Cells The Voltaic Cell
Mr. Shields Regents Chemistry U14 L03
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Half cell Nomenclature
Were now going to discuss some practical
applications of Redox reactions - Recall that
Redox reactions involve electron transfers -
Participating atoms either provide or accept
these e-
Example Zn0 ? Zn2 2e- Oxidation
half-cell Cu2 2e- ? Cu0 Reduction
half-cell
These reactions can also be written in
Shorthand nomenclature As Follows Zn0
Zn2 Cu2 Cu0
oxidation Reduction
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Electrochemical cells
The electrical nature of a Redox reaction allows
one the ability to construct several types of
electrical cells An apparatus that uses Redox
reactions to produce electrical Energy OR uses
electrical energy to cause a chemical reaction
is Known as an ELECTROCHEMICAL CELL
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Electrochemical cells
There are two types of electrochemical
cells - VOLTAIC (also known as GALVANIC)
CELLS - ELECTROLYTIC CELLS Devices that
convert CHEMICAL ENERGY into ELECTRICAL ENERGY
through a Spontaneous REDOX reaction are called
VOLTAIC CELLS They represent EXOTHERMIC reactions.
Note Well NYS Regents Refers to only the
VOLTAIC CELLS as an ELECTROCHEMICAL cell (even
though this is not quite accurate, remember it!)
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Electrochemical cells
In an ELECTROLYTIC CELL electrical energy is
provided to Force a NON-SPONTANEOUS Redox
reaction to happen. These Cells are ENDOTHERMIC
reactions. Electrolytic cells are typically
used to - plate metals on other metals -
obtain pure metal from its compounds - to
recharge batteries Well talk about GALVANIC
CELLS first and later well discuss ELECTROLYTIC
CELLS
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Remember We used Table J to determine if a Rxn
will occur spontaneously. A Rxn will be
spontaneous if the substance to be oxidized is
above the substance to be reduced. Example 1
Will the reaction Cu2 K ? Cu K be
spontaneous? Cu2 ? Cu (reduction) and K ? K
(oxidation) Since K is oxidized above Cu on
the table the Rxn is spontaneous.
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Example 2 Will the reaction Li Al ? Li
Al3 be spontaneous? Li ? Li
(reduction) Al ? Al3 (oxidation) Since Al is
oxidized but below Li on the table, the Rxn is
non-spontaneous.
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Electrochemical Cells
Most Voltaic cells share certain common
features. They all have - Two physically
separated half cells - A Cathode () in one cell
(Reduction occurs here) - An Anode (-) in the
other cell (Ox. Occurs here) - Two different
metal electrodes - A solution in each cell
containing a dissolved salt made from the metal
in the electrode - An electrical connection
between Anode and Cathode - A physical
connection between two separated cells
containing ions that can move freely into each
cell
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Memory Jogger for Galvanic cells
For Redox reactions youve remembered the phrase
OIL RIG For Voltaic cells there is another
phrase to help you remember how Voltaic cells are
constructed.
AN Ox ate a Red Pussy Cat
Anode Cathode
Oxidation Reduction
Negative Positive
A- oxidation C Reduction Fe ? Fe2 2e- Cu2
2e- ? Cu
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The basics of a Voltaic electrochemical cell
Anode Cathode
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Electrochemical Cells
Voltaic (galvanic) Cells generate usable
electricity Lets see how this happens.
Consider the following reaction
Fe CuCl2 ? FeCl2 Cu
This is a spontaneous Redox reaction. Fe is being
oxidized since its higher in Table J than Cu,
Cu will be reduced. What are the 2 half cell
reactions?
Fe ? Fe2 2e- oxidation half cell Cu2 2e- ?
Cu0 reduction half cell
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Electrochemical Cells
In this redox Reaction the Transfer of
e- Between Fe And Cu is Direct. In
Voltaic cells electron Transfer is indirect.
One way to do this is to separate the Fe and Cu
half cells and connect the metals by a wire.
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Electrochemical cells
If we do this, The reaction quickly Comes to a
Screeching halt. Why?
The build up of charge In both cells stops
the Reaction from Continuing. But Why
does Charge build up?
Fe2 goes into soln as Fe is oxidized releasing
2e- leaving behind pos. ions
The Cu2 consumes the 2 electrons leaving behind
an excess of Cl-. Rxn can no longer proceed.
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Salt Bridge

• We need a way to stop the build up of charge in
  each half cell.
• This is done using what is called a SALT BRIDGE
• - A salt bridge allows ions to flow from cell to
  cell but
• still maintains a separation
• between them
• It contains a salt in
• solution. The Cation () is
• be chosen such that it
• wont react with the metal
• Electrodes (Na salts are
• typical)

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Salt Bridge Ion diffusion
Pos. Ion Build Up
Neg. Ion Build Up
A membrane on each side of the salt bridge keeps
the NaNO3 Inside but allows ions to diffuse into
the half cells to keep them Electrically neutral
(i.e no build up of ions). In this case, for
each Cu reduced 2 Na ions must migrate into The
reduction half cell and at the same time 2 NO3-
ions Must migrate into the Oxidation cell (why?)
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Changing Concentration of Half cells
Notice that as the Cu2 is reduced the
concentration of Copper salt solution in that
half cell decreases. On the other hand, as Fe is
oxidized the concentration of Iron salt solution
in that half cell increases. WHY?
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Direction of electron Flow
In this voltaic cell e- transfer is from Fe to
Cu. Therefore the Fe Electrode is negative
(excess of e-) and Cu is (needs
electrons). Therefore electrons flow through the
wire from Fe to Cu.
But how do we know which metal is the One to
loose e- ?
Table J provides the Information needed. Metals
higher on the List are more easily Oxidized.
-

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Changing Electrode Mass
The electrode at which oxidation occurs loses
mass Fe ? Fe2 2e- The electrode at which
reduction occurs gains mass Cu2 2e- ? Cu
Since iron electrode is negative it is called
the And since Cu elec. is positive it is the
called the
ANODE (-)
-

CATHODE ()
Gains Mass
Looses Mass
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Voltage depends Upon the metal Electrodes used.

-
A standard Zn/Cu Electrochemical Cell produces
about 1.1V A standard Li/Ag Voltaic Cell produces
about 3.8V
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Equilibrium
As an electrochemical cell is used the voltage
continuously Decreases until the cell reaches
equilibrium.
When the oxidized electrode is completely used up
the cell voltage is zero (0) and the cell has
reached equilibrium since the concentrations are
no longer changing. The cell is now said to be
dead.
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PROBLEM
The following overall reaction occurs in a
galvanic cell. 4AgNO3(aq) Sn(s) ?
Sn(NO3)4(aq) 4Ag(s) Whats oxidized? Whats
Reduced? What are the balanced REDOX half cell
reactions? 4Ag 4e- ? 4Ag Reduction Half
cell Sn ? Sn4 4e- Oxidation Half
Cell What metal electrode gains mass? What metal
electrode loses mass? What electrode is negative,
which is positive? Which metal is the cathode,
Which is the anode? Which Soln Conc. Increases?
Sn Ag Ag Sn
Sn, Ag
Ag, Sn
Sn4
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Problem Draw and Fully label an Al Zn
Electrochemical cell. (Make sure you
Indicate the direction of e- and ion
flow use nitrate solutions ). Lastly, write
a balanced chemical equation that
describes this cell
Salt Bridge