Read Sections 8.3, and 8.4 before viewing the slide show.
Unit 29
Electrochemistry (Chapter 8)
•Description of an Electrochemical Cell (8.3)
•Electrochemistry Terminology (8.3)
•Electrochemistry as Applied to Batteries (8.3)
•Corrosion (8.4)
Electrochemistry (8.3)
•Electricity is due to the motion of electrons. Since oxidation-reduction
reactions involve an exchange of electrons, such reactions can be used to
generate electricity through applications such as batteries.
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Electrochemistry Cont. (8.3)
•In the reaction from the previous page, the
copper goes from being in the elemental form to
dissolving in solution as Cu2+ ions which cause
the blue color in solution.
•The Ag+ ions, previously dissolved in solution,
become elemental silver and form the “hangings”
seen in the second beaker.
•In equation form:
Cu (s) → Cu2+ (aq) + 2 eAg+ (aq) +
e-
→
Ag (s)
•An important aspect in understanding electrochemistry is to understand
how these two equations may be combined to form the overall reaction.
Electrochemistry Cont. (8.3)
•Each of the reactions below is called a half-reaction. One represents an
oxidation and the other a reduction.
Cu (s) →
Ag+ (aq) +
Cu2+ (aq)
e-
→
+
2 e-
Ag (s)
•Since the electrons donated by the copper are the ones accepted by the
silver, the number of electrons being accepted and donated must match.
•In order for the electrons to balance, each half-reaction is multiplied by an
integer as necessary to ensure that the number of electrons donated matches
those accepted. In this example, the Cu equation involves two electrons and
the Ag equation involves only one so multiplying the Ag equation by 2 will give
two electrons accepted to go along with the two electrons donated by copper
(continued on next page).
Electrochemistry Cont. (8.3)
•Multiplying the first equation by “1” and the second by “2” gives:
1 × (Cu (s) →
Cu2+ (aq)
2 × (Ag+ (aq) +
e-
+
→
2 e- )
Ag (s) )
•These simplify to:
Cu (s) →
Cu2+ (aq)
2 Ag+ (aq) +
2 e-
→
+
2 e-
2 Ag (s)
•Adding these two equations gives:
Cu (s) +
2 Ag+ (aq) → Cu2+ (aq) + 2 Ag (s)
Electrochemical Cells – Terminology (8.3)
Copper
Electrons
Silver
•Rather than carrying out the previous reaction
in one container, the two halves of the reaction
may be separated to allow the electrons to
transfer externally to the cell – see the figure to
the right.
•The semipermeable membrane allows the
nitrate ions to transfer through to the left
Anode
while the electrons transfer to the right
through the wire at the top.
•The copper metal connected to the wire is
called an electrode – specifically an anode
since that is where oxidation occurs.
•The silver metal is another electrode called
the cathode – the electrode at which reduction
occurs.
Cu2+
Ag+
SO42NO3
-
Semipermeable Membrane – only allows
solvent and nitrate ions to pass through.
Cathode
Electrochemical Cells – the Implementation (8.3)
•The figure below illustrates the construction of a dry cell typically used in
flashlights and other portable devices.
•A simplified version of the reaction that occurs is:
Zn + 2 MnO2 + H2O → Zn2+ + Mn2O3 + 2 OH•Alkaline cells replace use KOH in the paste – these
are typically more expensive but last longer.
•Can you tell which substance is oxidized – is it
zinc or manganese dioxide?
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The Lead Storage Battery (8.3)
•The lead storage battery is commonly found in cars and boats. It is a
rechargeable battery though it is quite heavy and involves corrosive materials.
•The typical 12-volt lead storage battery is made of six cells of two volts each.
•During the discharge of a lead storage battery (starting your car) the net
reaction is:
Pb + PbO2 + 2 H2SO4 →
2 PbSO4 + 2 H2O
•During the recharging, while the car is
running, the reverse reaction occurs
through the action of the car’s
alternator.
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Corrosion (8.4)
•Estimates are the corrosion in the US alone costs about $276 billion per year.
Approximately 20% of iron and steel production annually in the US is used to
replace corroded items.
•In the corrosion process, iron metal is initially oxidized to Fe2+ while oxygen in
the air is reduced to the hydroxide ion. This ultimately leads to iron (III)
hydroxide, which is the material commonly identified as rust.
•Electrons transferred in this process through the metal itself, but an electrolyte
is required to complete the circuit. Thus,
corrosion is more prevalent in northern
climates in which salt is used on the roads
and in areas near salt water.
•Often another metal that is more easily
oxidized is used as a “sacrifical” anode. Such
a material is destroyed preferentially to the
structural metal and is easily replaced. See
next slide.
Image from http://corrosionist.com
Corrosion Example
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Another Corrosion Example
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Sacrificial Anodes
•The small metal ingots (some highlighted in the image below) are called
“sacrificial” anodes.
•In a salt water environment, the “sacrificial” anodes will be destroyed
prior to the hull of the ship.
•Occasional replacement of the anodes is a relatively simple and
inexpensive task that does not affect the integrity of the hull.
Sacrificial Anodes
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