Read Sections 8.3, and 8.4 before viewing the slide show.
Slide 3
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)
Slide 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. Image from
http://www.vistatutor.com
Slide 5
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 Cu 2+ 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) Cu 2+ (aq) + 2 e - Ag + (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.
Slide 6
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) Cu 2+ (aq) + 2 e - Ag + (aq) + 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).
Slide 7
Electrochemistry Cont. (8.3) Multiplying the first equation by
1 and the second by 2 gives: 1 (Cu (s) Cu 2+ (aq) + 2 e - ) 2 (Ag +
(aq) + e - Ag (s) ) These simplify to: Cu (s) Cu 2+ (aq) + 2 e - 2
Ag + (aq) + 2 e - 2 Ag (s) Adding these two equations gives: Cu (s)
+ 2 Ag + (aq) Cu 2+ (aq) + 2 Ag (s)
Slide 8
Electrochemical Cells Terminology (8.3) Semipermeable Membrane
only allows solvent and nitrate ions to pass through. Anode Cathode
Copper Silver Cu 2+ SO 4 2- Ag + NO 3 - Electrons 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 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.
Slide 9
Image from http://lyrics.ashttp://lyrics.as 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 MnO 2 + H 2 O Zn 2+ + Mn 2 O 3 + 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?
Slide 10
Image from http://www.jamesglass.org 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 + PbO 2
+ 2 H 2 SO 4 2 PbSO 4 + 2 H 2 O During the recharging, while the
car is running, the reverse reaction occurs through the action of
the cars alternator.
Slide 11
Image from http://corrosionist.comhttp://corrosionist.com
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 Fe 2+
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.
Slide 12
Image from http://xtreme.hawaii.eduhttp://xtreme.hawaii.edu
Corrosion Example
Slide 13
Image from http://www.trekearth.comhttp://www.trekearth.com
Another Corrosion Example
Slide 14
Image from http://tis-gdv.dehttp://tis-gdv.de 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