Transcript 21.3 Electrolytic Cells

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Voltaic or
Galvanic Cells
D8 c34
Electrochemical
Cell
21.3 Electrolytic Cells >
Electrolytic vs. Voltaic Cells
Electrolytic vs. Voltaic Cells
How do voltaic and electrolytic
cells differ?
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21.3 Electrolytic Cells >
Electrolytic vs. Voltaic Cells
The process in which electrical energy
is used to bring about a chemical
change is called electrolysis.
• You are already
familiar with some
results of electrolysis,
such as gold-plated
jewelry, chromeplated automobile
parts, and silverplated dishes.
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The apparatus in which electrolysis is
carried out is an electrolytic cell.
21.3 Electrolytic Cells >
• An electrolytic cell uses electrical energy
(direct current) to make a non-spontaneous
redox reaction proceed to completion.
• An electrolytic cell is an
electrochemical cell used to cause
a chemical change through the
application of electrical energy.
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Electrolysis
Using electrical energy to produce chemical
change.
Sn2+(aq) + 2 Cl-(aq) ---> Sn(s) + Cl2(g)
SnCl2(aq)
Cl2
Sn
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21.3 Electrolytic Cells >
Electrolytic vs. Voltaic Cells
Voltaic Cell
Electrolytic Cell
Battery
e–
e–
e–
Anode
(oxidation)
Anode
(oxidation)
e–
Energy
Energy
Cathode
(reduction)
Cathode
(reduction)
In both voltaic and electrolytic cells,
electrons flow from the anode to the
cathode in the external circuit.
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21.3 Electrolytic Cells >
Electrolytic vs. Voltaic Cells
Voltaic Cell
Electrolytic Cell
Battery
e–
e–
e–
Anode
(oxidation)
Anode
(oxidation)
e–
Energy
Energy
Cathode
(reduction)
Cathode
(reduction)
The key difference between voltaic and electrolytic
cells is that in a voltaic cell, the flow of electrons is the
result of a spontaneous redox reaction, whereas in an
electrolytic cell, electrons are caused to flow by an
outside power source, such as a battery.
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21.3 Electrolytic Cells >
Electrolytic vs. Voltaic Cells
Voltaic Cell
Electrolytic Cell
Battery
e–
e–
e–
Anode
(oxidation)
Anode
(oxidation)
e–
Energy
Energy
Cathode
(reduction)
Cathode
(reduction)
• In a voltaic cell, the anode is the negative electrode and the
cathode is the positive electrode.
• In an electrolytic cell, the cathode is considered the
negative electrode.
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21.3 Electrolytic Cells >
In electrolysis, an electric current is
used to do which of the following?
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A.
Cause a chemical change
B.
Produce a battery
C.
Generate heat
D.
Run a motor
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21.3 Electrolytic Cells >
In electrolysis, an electric current is
used to do which of the following?
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A.
Cause a chemical change
B.
Produce a battery
C.
Generate heat
D.
Run a motor
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21.3 Electrolytic Cells >
Driving Nonspontaneous
Processes
Driving Nonspontaneous Processes
What are some applications that
use electrolytic cells?
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21.3 Electrolytic Cells >
Driving Nonspontaneous
Processes
Electrolysis of a solution or of a
melted, or molten, ionic compound
can result in the separation of
elements from compounds.
Electrolytic cells are also
commonly used in the plating,
purifying, and refining of metals.
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Electrolysis of
water
Anode (+)
2 H2O --->
O2(g) + 4 H+ +
4eCathode (-)
4 H2O + 4e- --->
2H2 + 4 OH-
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Electrolysis of Water
21.3 Electrolytic Cells >
The overall cell reaction is obtained by
adding the half-reactions (after doubling the
reduction half-reaction equation to balance
electrons).
Oxidation: 2H2O(l) → O2(g) + 4H+(aq) + 4e–
Reduction: 2 [2H2O(l) + 2e– → H2(g) + 2OH–(aq)]
6H2O(l) → 2H2(g) + O2(g) + 4H+(aq) + 4OH–(aq)
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21.3 Electrolytic Cells >
Driving Nonspontaneous
Processes
Electrolysis of Molten Sodium Chloride
The electrolytic cell in which this
commercial process is carried out is
called the Downs cell.
• The cell operates at
a temperature of
801°C so that the
sodium chloride is
maintained in the
molten state.
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21.3 Electrolytic Cells >
Electrolysis of Molten Sodium Chloride
Sodium and chlorine are produced
through the electrolysis of pure molten
sodium chloride, rather than an aqueous
solution of NaCl.
• Chlorine gas is produced at the anode.
• Molten sodium collects at the cathode.
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21.3 Electrolytic Cells >
Electrolysis of Molten Sodium Chloride
• The overall equation is the sum of the two
half-reactions:
Oxidation: 2Cl–(l) → Cl2(g) + 2e–
Reduction: 2Na+(l) + 2e– → 2Na(l)
2NaCl(l) → 2Na(l) + Cl2(g)
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21.3 Electrolytic Cells >
Electroplating
Electroplating is the deposition of a thin
layer of a metal on an object in an
electrolytic cell.
• An object that is to be silver-plated is made the
cathode in an electrolytic cell.
• The anode is the metallic silver that is to be
deposited.
• The electrolyte is a solution of a silver salt, such as
silver cyanide.
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21.3 Electrolytic Cells >
Electroplating
When direct current is applied, silver ions
move from the anode to the object to be
plated.
Reduction: Ag+(aq) + e– → Ag(s)
(at cathode)
This figure shows
statuettes that were
electroplated with
copper, nickel, and
24-carat gold.
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21.3 Electrolytic Cells >
Electrorefining
In the process of electrorefining, a piece of
impure metal is made into the anode of the
cell.
• It is oxidized to the cation and then
reduced to the pure metal at the
cathode.
• This technique is used to obtain
ultrapure silver, lead, and copper.
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21.3 Electrolytic Cells >
Electrowinning
In a process called electrowinning, impure
metals can be purified in electrolytic cells.
• The cations of molten salts or aqueous
solutions are reduced at the cathode to give
very pure metals.
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Producing Aluminum
2 Al2O3 + 3 C ---> 4 Al + 3 CO2
Charles Hall (1863-1914) developed
electrolysis process. Founded Alcoa.
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21.3 Electrolytic Cells >
Key Concepts
The key difference between voltaic and
electrolytic cells is that in a voltaic cell, the
flow of electrons is the result of a
spontaneous redox reaction, whereas in an
electrolytic cell, electrons are caused to
flow by an outside power source, such as a
battery.
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Electrolysis of a solution or of a melted, or
molten, ionic compound can result in the
separation of elements from compounds.
Electrolytic cells are also commonly used
in the plating, purifying, and refining of
metals.
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21.3 Electrolytic Cells >
Glossary Terms
• electrolysis: a process in which electrical
energy is used to bring about a chemical
change; the electrolysis of water produces
hydrogen and oxygen
• electrolytic cell: an electrochemical cell used
to cause a chemical change through the
application of electrical energy
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21.3 Electrolytic Cells >
BIG IDEA
Matter and Energy
• The two types of electrochemical cells
are voltaic cells and electrolytic cells.
• In an electrolytic cell, a nonspontaneous
redox reaction is driven by the
application of electrical energy.
• Electrolytic cells are used to produce
commercially important chemicals and
to plate, purify, and refine metals.
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Quantitative Aspects of
Electrochemistry
electrolysis of aqueous silver ion.
Ag+ (aq) + e- ---> Ag(s)
1 mol e- ---> 1 mol Ag
If measure the moles of e-, can know
the quantity of Ag formed.
But how to measure moles of e-?
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Quantitative Aspects of
Electrochemistry
• Measure the electrical current
charge passing
Current =
time
Units
coulombs
I (amps) =
seconds
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Quantitative Aspects of
Electrochemistry
I (amps) =
coulombs
charge passing
Current =
seconds
time
But how is charge related to moles of
electrons?
Ch arge on 1 mol e 
-19 C  
23 e - 
= 1.60 x 10
 6.02 x 10


e - 
mol 
=
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96,500 C/mol e-
=
1 Faraday
Michael Faraday
1791-1867
Originated the terms
anode, cathode, anion,
cation, electrode.
Discoverer of
• electrolysis
• magnetic props. of matter
• electromagnetic induction
• benzene and other organic
chemicals
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Quantitative Aspects of Electrochemistry
I (amps) =
32
coulombs
seconds
1.50 A flow thru a Ag+(aq) solution for 15.0 min.
What mass of Ag metal is deposited?
Solution: Calc. charge
Charge (C) = current (A) x time (t)
= (1.5 A)(15.0 min)(60 s/min)
= 1350 C
Quantitative Aspects of Electrochemistry
Charge = 1350 C
NEXT: Calculate moles of e- used
1350 C •
(c)
1 mol e  0.0140 mol e 96,500 C
Calc. quantity of Ag
1 mol Ag
0.0140 mol e - •
 0.0140 mol Ag or 1.51 g Ag
1 mol e -
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Pb(s) +
HSO4-(aq)
---> PbSO4(s) +
H+(aq)
+ 2e-
If a battery delivers 1.50 A, and you have 454 g
of Pb, how long will the battery last?
Solution
a)
454 g Pb = 2.19 mol Pb
b)
Calculate moles of e2 mol e 2.19 mol Pb •
= 4.38 mol e 1 mol Pb
c)
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Calculate charge
4.38 mol e- • 96,500 C/mol e- = 423,000 C
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NEXT:
Use Charge = 423,000 C
d)
Calculate time
Time (s) =
Charge (C)
I (amps)
423,000 C
Time (s) =
= 282,000 s
1.50 amp
About 78 hours