Transcript Electrochemical Cells

Electrochemical
Cells
https://egmanual.poly.edu/index.php?title=Lemon_Car
By: Maggie Dang
Background Information
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A chemical reaction that involves the transfer of electrons from one
substance to another is an oxidation-reduction (redox) reaction.
Experimentally, when copper wire is placed into a silver ion solution
copper atoms spontaneously lose electrons (copper atoms are
oxidized) to silver ions ( which are reduced). Silver ions migrate to the
copper atoms to pick up electrons and form silver atoms at the copper
metal/solution interface; the copper ions that form then move into the
solution away from the interface. The overall reaction that occurs at the
interface is
Cu (s) + Ag+ (aq) -> 2Ag(s) + Cu 2+ (aq)
This redox reaction can be divided into an oxidation and a reduction
half-reaction. Each half-reaction, called a redox couple, consists of the
reduced state and the oxidized state of the substance.
Cu(s) Cu2+ (aq) + 2e- (Oxidation half-reaction)
2Ag+ (aq) + 2e- 2Ag (s) (Reduction half-reaction)
http://members.chello.nl/r.kuijt/imag
es/en_oxidation_reduction.jpg
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A voltaic cell (galvanic cell) is designed to take advantage of this
spontaneous transfer of electrons. A voltaic cell separates the
copper metal from the silver ions and forces the electrons to pass
externally through a wire, an external circuit.
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The two redox couples are placed in separate compartments, called
half-cells. Each half-cell consists of an electrode, usually the metal
(reduced state) of the redox couple, and a solution containing the
corresponding cation (oxidized state) of the redox couple. The
electrodes of the half-cells are connected by a wire; this is where the
electrons flow, providing current for the external circuit.
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A salt bridge, which connects the two half-cells, completes the
construction of the voltaic cell (and the circuit). The salt bridge
permits limited movement of ions from one half-cell to the other, the
internal circuit, so that when the cell operates, electrical neutrality is
maintained in each half-cell.
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The electrode at which reduction occurs is called the
cathode; the electrode at which oxidation occurs is
called the anode.
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Because oxidation releases electrons to the electrode to
provide a current in the external circuit, the anode is
designated the negative electrode in a galvanic cell. The
reduction process draws electrons from the circuit and
supplies them to the ions in solution; the cathode is the
positive electrode.
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The cell potential of a galvanic cell is due to the
difference in tendencies of the two metals to oxidize or of
their ions to reduce.
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A measured reduction potential, the tendency
for a substance to gain electrons, is the value
used to identify the relative ease of reduction for
a half-reaction.
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A voltmeter, placed in the external circuit
between the two electrodes, measures the cell
potential, Eocell, a value that represents the
difference between the tendencies of the metal
ions in their respective half-cells to undergo
reduction.
Purpose
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To determine how closely the voltages found in
the experiment compare to the Ecello voltages
calculated using the standard reduction potential
chart
http://hsc.csu.edu.au/chemistry/opti
ons/industrial/2763/images/ch954_2
.gif
Materials
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Copper(II) nitrate Cu(NO3)2 1.0M
Iron(III) nitrate Fe(NO3)3, 1.0M
Lead(II) nitrate Pb(NO3)2 1.0M
Magnesium nitrate Mg (NO3)21.0M
Silver nitrate AgNO3 1.0M
Zinc nitrate Zn(NO3)2 1.0M
Sodium chloride NaCl 1.0M
2 beakers (50 mL )
Voltmeter
2 Alligator Clips
A salt bridge of 1.0M KNO3
Strips of Zn, Cu, and Mg
Iron nails
Lead Piping
Strips of filter paper
http://www2.gpmd.com/image/
b/bukm2040.jpg
Setup
http://mooni.fccj.org/~ethall/2046/ch18/zncu.gif
Copper
Procedures
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Prepare a test cell to measure the voltage of the
copper and zinc half-cells.
Put approximately 2 mL of 1.0 M Cu(NO3)2 in
one beaker and 2 mL of 1.0 M Zn(NO3)2 in the
other beaker.
Polish small strips of zinc and copper metal, and
place the metal in the appropriate beaker
containing the solution of the ions of that metal.
http://www.humboldtmfg.co
m/images/products/H4911beakers.jpg
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http://www.svmetal.com/images/
copper_image.jpg
ZINC
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http://www.united
nuclear.com/zincs
tick.jpg
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Take a small strip of filter paper that has been
soaked in KNO3 solution, and drape it across
the wells so that one end dips in the solution in
each well. This will act as the salt bridge.Use a
voltmeter to measure the voltages between the
two half-cells.
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Connect the meter so the voltage reading is
positive. When the voltmeter reads a positive
voltage, the electrode connected to the positive
terminal is the cathode and is undergoing
reduction, while oxidation is occurring at the
electrode connected to the negative terminal,
the anode.
http://www.chem.fsu.edu/che
mlab/chm1046lmanual/electr
ochem/boilkno3.JPG
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Record data.
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Rinse out the beakers and measure the voltages
of the other electrodes the same way with zinc
since it is designated to be the standard
electrode.
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The electrodes to be tested are : Ag, Ag+; Cu,
Cu+; Fe, Fe3+; Mg,Mg2+; Pb, Pb2+ ; Zn, Zn2+
Voltmeter
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Record data.
http://car-batteries.org.uk/wpcontent/uploads/2009/03/voltmeter-carbattery-hull.jpg
Voltage of each half-cell versus the zinc electrode
Voltage
Anode
Cathode
Zn versus Ag 1.40
Zn
Ag
Zn versus Cu .99
Zn
Cu
Zn versus Fe .55
Zn
Fe
Zn versus Mg .60
Mg
Zn
Zn versus Pb .48
Zn
Pb
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Write reduction equations for each metal
ion.
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Record the standard potentials for all of
the electrodes using the standard
reduction potentials chart, and calculate
the potential energy of the entire cell by
using the equation
Eocell=Eored (cathode) – Eored (anode)
Calculations
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Standard Electrode Zinc with Eo= -.76
Ag+ + e-  Ag Eo=.80
.80 - -.76=1.56
Cu2+2e-Cu Eo=.34
.34 - -.76=1.1
Fe+3 +3- Fe Eo= -.44
-.44 - -.76= .32
Mg2+ 2e-  Mg Eo= -.76
-.76 - -2.37 =1.61
Pb2+ 2e- Pb Eo= -.13
-.13 - -.76=.63
Zn2+ 2e- Zn Eo= -.76
-.76 - -.76 = 0
Reduction Equations for Each Ion Arranged in Decreasing Order
of Potential
Reduction
Reaction
Voltages
using Zinc
as the
Standard,
EZn
Eored Voltages Using Standard Reduction Potential Chart
Ag+ + e- 
Ag
1.40
1.56
Cu2+ + 2eCu
.99
1.10
Fe3+ + 3eFe
.55
.32
Mg2+ + 2e Mg
.60
1.61
Pb2+ + 2ePb
.48
.63
Zn2+ + 2eZn
0.00
0.00
Conclusion
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Comparing to the Eored voltages calculated from
using the standard reduction potential chart, the
voltages measured using Zinc as the standard,
EZn, in the experiment are relatively close.
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Therefore, this is an accurate experiment that
can be used to measure potential energies
among different types of electrodes.
Sources
Vonderbrink, Sally Ann. Laboratory
Experiments for Advanced Placement
Chemistry Student Edition. Flinn Scientific,
Inc. Batavia, IL. 1995
 Beran, J.A. Laboratory Manual for
Principles of General Chemistry. John
Wiley & Sons, Inc. Hoboken, NJ. 2004.
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