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ELECTROCHEMISTRY
Chapter 21
Section 21.1 Electrochemical Cells
• Objectives
– Interpret an activity series and identify the
elements that are most easily oxidized and
those that are least easily oxidized
– Name the type of reactions involved in
electrochemical processes
– Describe how a voltaic cell produces electrical
energy
• Fireflies
– Glow to attract mates
• Anglerfish
– Emit light to attract prey
• Squid, jellyfish, bacteria, and
shrimp
– Luminous
How?
• Redox reactions
• Transfer of
electrons in a redox
reaction provides
energy
Electrochemical Processes
• Chemical processes either release or absorb
energy
– Energy is sometimes in the form of electricity
• Electron transfer reactions or redox reactions
– result in the generation of an electric current
spontaneously or
– are caused by imposing an electric current
nonspontaneously
– The field of chemistry that deals with these two
situations is called electrochemistry
Spontaneous Redox Reactions
How do we know this reaction between zinc and
copper is spontaneous?
– Table J – Activity Series
• Zinc is higher on the list than copper
–For any two metals on table J, the more
active metal is more readily oxidized
Which are spontaneous & if spontaneous
then Identify Reaction Products ?
• Li + AlCl3 
•
•
•
•
•
•
Spontaneous (Li above Al)
Cs + CuCl2  Spontaneous (Cs / Cu)
I2 + NaCl  Nonspontaneous (Cl / I)
Spontaneous (Cl / Br)
Cl2 + KBr 
Fe + CaBr2  Nonspontaneous (Ca / Fe)
Nonspontaneous (Sr / Mg)
Mg + Sr(NO3)2 
F2 + MgCl2  Spontaneous (F / Cl)
Electrochemistry
• Many applications
– Flashlights
– Automobile batteries
– Manufacture of sodium and aluminum metal
– Silver plating of table ware and jewelry
– Biological systems
• Carrying impulses
Electrochemical Cells
• An electrochemical cell is any device that converts
chemical energy into electrical energy or electrical
energy into chemical energy
– Redox reactions occur in all electrochemical cells
• A battery functions by transferring electrons
through an external wire from the reducing agent to
the oxidizing agent.
Overview of Electrochemistry
Two kinds of electrochemical cells
(kind of opposites):
1. Voltaic
• Use a spontaneous reaction
to produce a flow of
electrons (electricity) exothermic.
2. Electrolytic
• Use a flow of electrons
(electricity) to force a
nonspontaneous reaction to
occur - endothermic.
Batteries are voltaic
cells
Animation of Voltaic Cell
http://www.chem.iastate.edu/group/Greenbo
we/sections/projectfolder/flashfiles/electroCh
em/volticCell.html
Voltaic Cells
• Spontaneous redox reactions
• Converts chemical energy into electrical
energy
– Electrical energy is produced in a voltaic cell by
spontaneous redox reactions within the cell
• Parts of a Voltaic Cell
– 2 half cells
– 2 electrodes (an anode and a cathode)
– Aqueous solutions
– Wire
– Salt bridge
Voltaic Cell
Anode
Cathode
Aqueous solution must contain
ions of same metal as electrode:
here ions = Zn2+ ions.
Aqueous solution must contain
ions of same metal as electrode:
here ions = Cu2+ ions.
Solution might be
Zn(NO3)3(aq) or ZnSO4(aq)
Solution might be
Cu(NO3)3(aq) or CuSO4(aq)
Voltaic Cell
Wire: connects the electrodes, carries electrons (electric current)
Anode
Cathode
Salt bridge: Allows ions to pass from one cell to another
but prevents solutions from mixing completely
Anode and Cathode
• Electrodes (the metals)
– Anode
• Electrode at which the oxidation occurs
– Electrons are produced at the anode and it is labeled
the negative electrode
– Cathode
• Electrode at which reduction occurs
– Electrons are consumed at the cathode as a result the
cathode is labeled the positive electrode
– Neither electrode is really charged
• All parts remain neutral
• Moving electrons and ions balance any charge that might
build up
What happens at the electrodes?
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Anode:
Zinc metal slowly dissolves
(looses mass)
Oxidation: Zn0  Zn2+ + 2eCathode:
Copper atoms are deposited
as metallic copper on top of
zinc (gaining mass)
Reduction: Cu2+ +2e-  Cu0
An Ox Ate a Red Cat
• Anode – Oxidation
– The anode = location for the
oxidation half-reaction.
• Reduction – Cathode
– The cathode = location for the
reduction half-reaction.
How do we determine which electrode is the
anode and which electrode is the cathode?
• Remember…
– Table J
• The more active metal is oxidized and is therefore
the anode
• Anode = oxidation = electron donor
– Higher metal on table J
• Cathode = reduction = electron acceptor
– Lower metal on table J
Zn is above Cu on Table J
Zn is the anode
Cu is the cathode
Zn2+
Electrons flow from zinc to copper (through wire)
Positive ions (Zn2+)flow from zinc to copper (through salt bridge)
Negative ions (SO42-)flow from copper to zinc (through salt bridge)
e-
e-
e-
e-
Anode
e-
Cathode
e-
Complete Electrochemical Cell
Zn --> Zn2+ + 2e-
Cu2+ + 2e- --> Cu
Oxidation
Anode
Negative
AN OX
Reduction
Cathode
Positive
<-- Anions
Cations -->
RED CAT
How a Voltaic Cell Works …
•
The electrochemical process that occurs in a Zn-Cu voltaic cell can best
be described in steps, but the steps actually occur at the same time
1.
Electrons are produced at zinc anode according the
oxidation half reaction: Zn(s)  Zn2+(aq) + 2e-. Zinc is the
anode and the negative electrode because it is oxidized
Electrons leave zinc anode and pass through the external
circuit (wire) to the copper rod
Electrons enter the copper rod and interact with copper ions
in solution. The following reduction half reaction occurs
Cu2+(aq) + 2e-  Cu(s). Copper ions are reduced. The
copper is the cathode and the positive electrode because it
is reduced
To complete the circuit, both the positive and negative ions
move through the aqueous solution via the salt bridge
2.
3.
4.
Construct a Voltaic Cell with Al & Pb
1. Use Table J to identify anode & cathode.
2. Draw Cell, put in electrodes & solutions using
nitrate as the negative ion
3. Label:
a. anode
b. cathode
c. direction of electron flow in wire
d. direction of positive ion flow in salt bridge
e. positive electrode
f. negative electrode.
4. Write out half reactions
5. Write the overall reaction
Electron flow 
wire
Positive ion flow 
Al =
anode
Pb =
cathode
Salt bridge

Al+3
& NO3
-1
Pb+2 & NO3-1
What are the half-reactions?
Al 
Al+3
+
3e-
Al metal is the electrode – it is
dissolving. Al+3 ions go into the
solution. (loosing mass)
Pb+2 + 2e-  Pb
Pb+2 ions are in the solution. They
pick up 2 electrons at the surface of
the Pb electrode & plate out.
(gaining mass)
Al is higher
on Table J
= anode
= oxidation
Pb is lower
on Table J
= cathode
= reduction
Overall Reaction
2(Al  Al+3 + 3e-)
3(Pb+2 + 2e-  Pb)
____________________________
2Al + 3Pb+2  2Al+3 + 3Pb
2Al + 3Pb+2  2Al+3 + 3Pb
• Which electrode is losing mass? Al
• Which electrode is gaining mass? Pb
• What’s happening to the [Al+3] in solution?
Increasing
• What’s happening to the [Pb+2]in solution?
Decreasing
Animation of Voltaic Cell
Let’s see the animation again:
http://www.chem.iastate.edu/group/Greenbo
we/sections/projectfolder/flashfiles/electroCh
em/volticCell.html
Videos
• http://www.youtube.com/watch?v=kfgtU9D
DvdY&feature=channel (3:29 min)
21.3
Section 21.3 – Electrolytic Cells
• Objectives
– Distinguish between electrolytic and voltaic cells
– Describe the chemical changes that take place
during electrolysis
– Name three ways that electrolysis is used in metal
processing
Electrolytic vs. Voltaic Cells
• Voltaic cells convert chemical energy to electrical
energy during a spontaneous redox reaction
• Electrolytic cells use an electric current to make a
nonspontaneous redox reaction go forward
– This process of using electrical energy to bring about a
chemical change is called electrolysis
• Applications of electrolysis include:
– Silver plating dishes and utensils
– Gold-plating jewelry
– Chrome plating automobile parts
Voltaic vs Electrolytic Cells
– Student version
Answer Key
Voltaic & Electrolytic Cells
What’s the same?
• Anode is always the site of oxidation
– AN OX
• Cathode is always the site of reduction
– RED CAT
• Electrons always flow from anode to
cathode
Voltaic & Electrolytic Cells
What’s different?
• Presence of a battery
– Voltaic – no battery
– Electrolytic – has battery
• Which metal is the anode & which is the cathode
– Voltaic – anode is more active metal  spontaneous
– Electrolytic – anode is less active metal  non-spontaneous
• Charge on anode & charge on cathode
– Voltaic – anode is +
– Electrolytic – cathode is -
Electroplating
• Example of Electrolysis
• One metal is plated onto another
• Used to protect the surface of the base metal from corrosion
• Or to make objects, such as tableware and jewelry more attractive
• An electric current is used to produce a
chemical reaction
– The object to be plated is the cathode
(negative)
– The metal used for platting is the anode
(positive)
Electroplating
e-
Example
– copper plating a key
• The key is the cathode
• Copper is the anode
Cu Anode
oxidation
Cu dissolves
Cu2+ ions in solution
combine with excess eon key to form Cu
coating
Electroplating
• Usually, the object to be electroplated, such as a spoon, is cast of an
inexpensive metal. It is then coated with a thin layer of a more attractive,
corrosion-resistant, and expensive metal, such as silver or gold.
Oxidation…
so is Ag the
anode
anode or
cathode? Oxidation
Look at the direction
of e- flow, what does
that tell us about
where oxidation
occurs?
Ag0  Ag+ + 1eCathode
Ag+ + e-  Ag0
Silver coats spoon
Electrolysis of Molten Sodium Chloride
• Our next example is the electrolysis of
molten sodium chloride
– This is the only way to produce sodium metal
• Why do you think we need to melt the
sodium chloride first?
– So the ions are mobile
– when you hear molten NaCl – think of ions
floating around in liquid
Let’s examine the electrolytic cell for molten NaCl.
Which way do the electrons flow through the wire?
Electrons always flow from anode to cathode…which is the anode?
What is the charge on the anode? What is the charge on the cathode?
What electrode are the Cl- attracted to? The Na+?
+
-
-
+
Inert electrode
e-
ClNa+
Anode
Ox: 2Cl-  Cl2 + 2e-
Na+
Cl-
Inert electrode
e-
Cathode
Red: Na+ + 1e-  Na
 Chlorine ions combine to form chlorine gas
 Remember this is molten NaCl, when all the Cl escapes as gas, we
are left with Na
e-
+
-
e-
Inert electrode
ClCl2(g)
eCl-
Anode
Ox: 2Cl-  Cl2 + 2e-
0
Na
Na+
Na0+
e-
Inert electrode
-
+
Cathode
Red: Na+ + 1e-  Na
Electrolysis of Molten NaCl
Observe the reactions at the electrodes
-
battery
+
Cl2 (g) escapes
Na (l)
Clelectrode
half-cell
(-)
ClCl-
Na+ + e-  Na
NaCl (l)
Na+
Na+
(+)
Na+
electrode
half-cell
2Cl-  Cl2 + 2e-
Predicting Spontaneous Redox Reactions
In the 2 beakers a strip of Cu
was placed in a solution
Zn(NO3)2 or AgNO3.
Which beaker had the Zn(NO3)2
& which had AgNO3?
Beaker (a) was the solution
of AgNO3 and beaker (b) was
the solution of Zn(NO3)2
Because there is no outside source of electricity the only reaction that
can occur is a spontaneous one – Cu is above Ag therefore
spontaneous, Cu is below Zn, therefore a nonspontaneous reaction
and would require an outside source of energy to force the reaction to
occur
End of Ch 21
Corrosion Prevention
Zinc is more easily oxidized than iron, therefore the
zinc will be oxidized as opposed to the iron prevent
corrosion of the iron metal
Application: Corrosion
Charging a Battery
When you charge a battery, you are forcing
the electrons backwards (from the + to the -).
To do this, you will need a higher voltage
backwards than forwards. This is why the
ammeter in your car often goes slightly higher
while your battery is charging, and then
returns to normal.
In your car, the battery charger is
called an alternator. If you have a
dead battery, it could be the
battery needs to be replaced OR
the alternator is not charging the
battery properly.
Predicting Redox Reactions
A + BX  B + AX
A & B are metals. If metal A is above
metal B in Table J, the reaction is
spontaneous.
X + AY  Y + AX
• X & Y are nonmetals. If nonmetal X is
above nonmetal Y in Table J, the
reaction is spontaneous.
Electrolysis of Molten NaCl
At the microscopic level
-
e-
battery
+
NaCl (l)
cations
migrate
toward
(-)
electrode
Cl-
(-)
cathode
Na+
+
e-
Na+
ClCl-
 Na
Na+
e(+)
Na+
anions
migrate
toward
(+)
electrode
anode
2Cl-  Cl2 + 2e-