Electrochemistry: Voltaic Cells

Experiment 11ELECTROCHEMISTRY: VOLTAIC CELLS
1
Electrochemistry: Voltaic Cells
OBJECTIVES
• Prepare a Cu-Pb voltaic cell and measure its potential.
• Test two voltaic cells that use unknown metal electrodes and identify the metals.
• Construct a Cu-Zn soil battery and measure its voltage.
• Connect several of the soil battery cells to obtain higher voltages and light different color
LEDs.
LECTURE TOPIC REFERENCES
Review the following before performing the experiment:
Chapter Title: Electrochemistry
Section Title: Balancing Oxidation-Reduction Equations; Voltaic (or Galvanic Cells; Standard Electrode
Potential; Batteries: Using Chemistry to Generate Electricity
Tro, N. (2017). Chemistry: A Molecular Approach. 4th Edition. Boston: Pearson. (Or latest edition)
CONCEPTS
Electrochemistry is the study of chemical reactions that produce electricity. Such chemical
reactions involve electron transfer, and are therefore oxidation-reduction (redox) reactions. If a
redox reaction occurs spontaneously in a cell, it is called a voltaic (or galvanic) cell. In this
experiment, you will study the characteristics and operation of several voltaic cells. You will then
construct a battery, which is a practical application of voltaic cells.
Redox Reaction
In a redox reaction two half reactions occur; one reactant gives up electrons (oxidation at the
anode) and another reactant gains electrons (reduction at the cathode). For example:
Oxidation half-reaction (anode):
2+
–
Zn(s) → Zn (aq) + 2e
(Equation 1)
In this half reaction, the oxidation number of Zn(s) is 0 and increases to +2.
2+
–
Reduction half-reaction (cathode): Cu (aq) + 2e → Cu(s)
2+
E0 = +0.34V
In this half reaction, the oxidation number of Cu (aq) is 2+ and is reduced to 0.
(Equation 2)
2
However, no half reaction can occur by itself and the two half reaction are combined to form the
redox reaction. Thus, for the above example half-reactions, the overall redox reaction can be
obtained by adding equations 1 and 2,
2+
Zn(s) → Zn (aq) + 2e
2+
–
–
Cu (aq) + 2e → Cu(s)
——————————————-2+
2+
Zn(s) + Cu (aq) → Zn (aq) + Cu(s)
(Equation 3)
The electrons for the reaction are not shown because they are neither reactants nor products, but
2+
have simply been transferred from one species to another (from Zn to Cu in this case).
Standard Reduction Potential
We cannot measure the absolute tendency of a half-reaction; we can measure it only relative to
another half-reaction. We have selected the reduction of H+ to H2 under standard conditions (25oC
and 1 atm. H2 gas pressure), as a standard half-reaction, and assign this a potential of 0.00V. Any
half-reaction can be expressed as a reduction, and thus the oxidation reaction of Zn, equation (1),
can be reversed to be presented as a reduction reaction.
2+
–
E0 = -0.76V
Zn (aq) + 2e → Zn(s)
(Equation 4)
A measure of the tendency for a reduction to occur is its reduction potential, E, measured in units
of volts. At standard conditions, 25 °C and concentrations of 1.0 M for the aqueous ions, the
measured voltage of the reduction half-reaction is defined as the standard reduction potential, E°.
Standard reduction potentials have been measured for most half-reactions and they are listed in
Table 1 in Attachment A. The more positive (or less negative) the reduction potential, the greater is
2+
2+
the tendency for the reduction to occur. So, Cu has a greater tendency to be reduced than Zn .
Furthermore, Zn has a greater tendency to be oxidized than Cu.
Voltaic Cells
The voltaic (galvanic) cell is an electrochemical cell that produces electric current from a
spontaneous redox chemical reaction. A voltaic cell based on the Equation 3 redox reaction is
shown in the Figure 1. The voltaic cell consists of; an anode: an electrode where oxidation always
occurs, it is the more negatively charged electrode and denoted by the negative sign (-), a cathode :
an electrode where reduction always occurs, it is the more positively charge electrode and denoted
by a positive sign (+), a salt bridge containing the electrolyte that allows the ‘current’ of ions to flow
from the anode and the cathode, and an external conducting wire connecting the anode and
cathode to produce an electric current. The difference in the potential energy between the anode
and cathode of voltaic cell is the cell potential and measured as volts (V), while the number of
electrons flowing from the cell per second is measured by the unit ampere (A).
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Figure 1
Cu-Zn Voltaic Cell
In this experiment, you will prepare a variety of semi-microscale voltaic cells in a 12-well test plate
(Figure 2). A voltaic cell is constructed by using two metal electrodes and solutions of their
respective salts (the electrolyte component of the cell) with known molar concentrations. In Parts I
and II of this experiment, you will use a Voltage Probe to measure the potential of a voltaic cell with
copper and lead electrodes. You will then test two voltaic cells that have unknown metal electrodes
and, through careful measurements of the cell potentials, identify the unknown metals.
Figure 2 Well Test Plate Voltaic Cell
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The voltage of any voltaic cell is the difference of the reduction half-reaction occurring at the
cathode (+) minus the reduction half-reaction occurring at the anode (-).
Ecell = Ered(+) – Ered(-)
(Equation 5)
Thus, the voltaic cell potential (Ecell) can be calculated by considering the standard reduction
potential of the individual half-reactions of the redox reaction listed in Table 1 (Attachment A).
For example, considering the redox reaction from Equation 3,
2+
2+
Zn(s) + Cu (aq) → Zn (aq) + Cu(s)
The voltaic cell voltage is obtained from the reduction half-reaction shown in Equations 2 and 4 and
from Table 1. Using Equation 5, the voltaic cell potential (Ecell) is
2+
–
2+
–
Cu (aq) + 2e → Cu(s)
Zn (aq) + 2e → Zn(s)
E0 = +0.34V
E0 = -0.76V
Ecell = 0.34V- (-0.76) = +1.10 V
(Equation 6)
In a similar manner, if the voltaic cell voltage can be measured and one half-reaction is known, the
other reaction can be determined. For example, if the cell voltage is measured to be +0.53 V, and
the redox reaction consists of a zinc electrode and another metal, we would insert the value of cell
voltage and the reduction half-reaction of Zn in Equation 5 to obtain the reduction half-reaction of
the other reaction.
Ecell = Ered(+) – Ered(-)
0.53V = Ered(+) –(-0.76V)
Ered(+) = 0.53V – 0.76 V = -0.23V
Table 1 in the attachment A shows that the half-reaction potential is -0.23V if the reaction is
2+
–
Ni (aq) + 2e → Ni(s)
Thus the unknown electrode is Ni.
E0 = -0.23V
(Equation 7)
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Batteries
Batteries are the most familiar devices for converting between electrical and chemical energies. A
battery is a self-contained power source that consists of one or more voltaic cells. A primary
battery, such as a 1.5V AA alkaline cell, is used once and discarded as the electrode materials are
irreversibly changed during use. A secondary battery, such as a Li-ion battery in your cell phone,
can be discharged and charged several times. The change in electrode composition during
discharge is restored to its original composition during recharge (reverse current).
A battery may consist of a single voltaic cell, such as 1.5V AA alkaline cell. Many voltaic cells may be
connected in series to obtain higher voltages to power electronic devices. For example, one inserts
3 alkaline cells to light a flashlight, or a 12V automotive starter battery consists of six 2V lead acid
cells.
Alessandro Volta built the first electrochemical battery in 1799, called the ‘voltaic pile’. It consisted
of copper and zinc plates separated by brine soaked paper that acted as the electrolyte. In this
experiment we will construct a soil battery consisting of Cu and Zn electrodes and wet soil as the
electrolyte. We will then connect several of these voltaic cells to obtain higher voltages to light
different color light emitting diodes (LED).
Light Emitting Diode (LED)
A LED (Light Emitting Diode) is a semiconducting device that produces light when a current flows
through it. The device consists of a semiconducting structure of p-n (positive-negative) junction
that will allow current flow from the p-side (anode) to the n-side (cathode), but not in the reverse
direction. When charge carriers (electrons and holes) meet at the junction they fall to a lower
energy and release energy as photons (light), Figure 3.
The specific wavelength or color emitted by the LED depends on the materials used to make the
diode as shown in the Table 2 (Attachment A) and Figure 4 below. Additionally, the intensity of
light is proportional to the magnitude of the current flowing through the LED. Thus, to obtain a
given intensity of light at each color, each material requires a different applied forward operating
voltage (VF).
Figure 3
Light Emitting Diode (LED)
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Figure 4 – LED Characteristics
MATERIALS
Vernier computer interface (LabQuest 2)
Computer
Logger Pro
Voltage Probe
12-well test plate
Cotton string (about 2.5 to 3-inch length)
Cu and Pb electrodes
2 unknown electrodes, labeled X and Y
150 mL beaker
0.10 M copper (II) nitrate, Cu(NO3)2, solution
0.10 M lead (II) nitrate, Pb(NO3)2, solution
1 M potassium nitrate, KNO3, solution
0.10 M X(NO3)n solution
0.10 M Y(NO3)m solution
Sand paper
Plastic Beral pipets
1 ice-cube tray
4 Cu electrodes 2”x1”
4 Zn electrodes 2”x1”
2 red alligator clip leads
2 black alligator clip leads
300 mL of Soil
Soil water
LED lights- 1 each of red, yellow, white, green and blue color.
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PRE-LAB ASSIGNMENT
Use the table of standard reduction potentials (Table 1 in Attachment A), or another approved
reference, to complete the Table 3 below, by constructing a voltaic cell of each electrode couple.
For the example provided below, see calculations in Equation 6 in the Concept section.
Table 3 Voltaic Cell Potentials (Ecell) of Some Electrode Couples
Electrodes
Half-Reactions
E°
E°cell
Zn
Zn(s)→Zn2+ + 2e–
+0.76 V
+1.10 V
Cu
Cu2+ + 2e– → Cu(s)
+0.34 V
Cu
Pb
Pb
Ag
Pb
Mg
Pb
Zn
PROCEDURE
Part I
Determine the E° for a Cu-Pb Voltaic Cell
1. Obtain and wear goggles.
2. Use a 12-well test plate as your voltaic cell. Use Beral pipets to transfer small amounts (1 mL or
20 drops) of 0.10 M Cu(NO3)2 and 0.10 M Pb(NO3)2 solution to two neighboring wells in the test
plate. See Figure 5 for proper distances of voltaic cells to avoid cross contamination.
Figure 5 Suggested Locations of Voltaic Cells in a 12-Well Test Plate
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WARNING: Copper (II) nitrate, 0.2 M, Cu(NO3)2: Causes skin and serious eye irritation.
DANGER: Lead nitrate solution, Pb(NO3)2: Do not eat or drink when using this product—harmful
if swallowed or inhaled. Avoid breathing dust and fumes. May damage fertility or the unborn
child. Inorganic lead compounds are probable human carcinogens. Do not handle until all safety
precautions have been understood.
3. Obtain one Cu and one Pb metal strip to act as electrodes.
DANGER: Solid lead, Pb: Do not eat or drink when using this product—harmful if swallowed or
inhaled. Avoid breathing dust and fumes. May damage fertility or the unborn child. Inorganic
lead compounds are probable human carcinogens. Do not handle until all safety precautions
have been understood. Polish each strip with steel wool. Place the Cu strip in the well of
Cu(NO3)2 solution and place the Pb strip in the well of Pb(NO)3 solution. These are the half cells
of your Cu-Pb voltaic cell.
4. Make a salt bridge by soaking a short length of cotton string in a beaker that contains a small
amount (about 1 mL or 20 drops) of 1 M KNO3 solution (If presented as soaked strings to you,
use the forceps to transfer the wet strings in the container provided). Connect the Cu and Pb
half cells with the string.
WARNING: Solid potassium nitrate, KNO3: May intensify fire—oxidizer. Keep away from heat,
sparks, open flames, and hot surfaces. May be harmful if swallowed.
5. Connect a Voltage Probe to Channel 1 of the Vernier computer interface (LabQuest 2). Connect
the interface to the computer with the proper cable (optional, because you can measure the
voltage directly from the LabQuest 2).
6. Start the Logger Pro program on your computer. Open the file “20 Electrochemistry” from the
Advanced Chemistry with Vernier folder.
7. Measure the potential of the Cu-Pb voltaic cell. Complete the steps quickly to get the best data.
a. Click
to start data collection. Note: When the voltage measurement leads are not in
contact with a cell (or each other), a meaningless default voltage may be displayed. If you
touch the two leads together, the voltage will drop to about 0.00 V.
b. Connect the leads from the Voltage Probe to the Cu and Pb electrodes to get a positive
potential reading. Click
immediately after making the connection with the Voltage
Probe.
c. Remove both electrodes from the solutions. Clean and polish each electrode using a small
piece of sand paper (make sure to remove any sand particles left on the electrode).
d. Put the Cu and Pb electrodes back in place to set up the voltaic cell. Connect the Voltage
Probe to the electrodes, as before. Click
immediately after making the connection
with the Voltage Probe.
e. Remove the electrodes. Clean and polish each electrode again (as in step 7.c).
f. Set up the voltaic cell a third, and final, time. Click
immediately after making the
connection with the Voltage Probe. Click
to end the data collection.
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g. Click Statistics, . Record the mean in the Data and Calculations section as the average
potential. Save the graph for submission in your lab report, and close the statistics box on
the graph screen by clicking the X in the corner of the box.
h.
As an alternative, you can record the highest potential of the first determination instead
of getting an average.
Note: See sample data below.
Part II
Determine the Eo for Two Voltaic Cells Using Pb and Unknown Metals
8. Obtain a small amount of the unknown electrolyte solution labeled “0.10 M X” and the
corresponding metal strip, X.
9. Use a Beral pipet to transfer a small amount (about 1 mL or 20 drops) of 0.10 M X solution to a
well adjacent to the 0.10 M Pb(NO3)2 solution in the test plate (leave one empty well to avoid
cross contamination).
DANGER: X solution, X(NO3)n: Can intensify fire (an oxidizing agent), and can cause eye and skin
irritation. Y solution, Y(NO3)m: A suspected carcinogen.
10. Using a pair of forceps, obtain a short length of cotton string soaking in the beaker of 1 M KNO3
solution and place in a small plastic weighing boat. Connect the X and Pb half cells with this
KNO3 salt bridge.
11. Measure the potential of the X-Pb voltaic cell. Complete this step quickly.
a. Click
to start data collection.
b. Connect the leads from the Voltage Probe to the X and Pb electrodes to get a positive
potential reading. Click
immediately after making the connection with the Voltage
Probe.
c. Remove both electrodes from the solutions. Clean and polish each electrode.
d. Set up the voltaic cell again. Connect the Voltage Probe as before. Click
after making the connection with the Voltage Probe.
immediately
e. Remove the electrodes. Clean and polish each electrode again (as in Step 7.c).
f. Test the voltaic cell a third time. Click
the Voltage Probe.
g. Click
immediately after making the connection with
to end data collection.
h. Click Statistics, . Record the mean in the Data and Calculations section as the average
potential. Save the graph for submission in your lab report, and close the statistics box on the
graph screen by clicking the X in the corner of the box. As an alternative, you can record the
highest potential of the first determination instead of getting an average.
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i. As an alternative, you can record the highest potential of the first determination instead of
getting an average.
12. Repeat Steps 8–11 using the unknown and its corresponding electrolyte solution labeled “Y”.
Note: See sample data below.
EXPERIMENT VIDEO
The instructor will explain this part of the experiment using Figures 2 and 5 above instead of
showing a video.
Part I and II SAMPLE DATA
Table Sample Data 1
Cell Potential Readings
Metal Electrodes
Cu/Pb
X/Pb
Y/Pb
Cell Potential (V)
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Part III Making of Soil Battery and Lighting LED Lamps
1. Prepare the soil electrolyte. Obtain about 300 mL of soil sample in a 1000 mL beaker
and add about 100 mL of “Soil Water.” (Do not attempt to add DI water to the
mixture).
2. Distribute the soil electrolyte to 4 neighboring cells in the ice-cube tray in a loose
manner. The soil will fill the cells but not overflow (about ¾ full, and soil should not
touch the nearby cells).
3. Starting from the cell on the left, insert one Cu (red color in the images below) and one
Zn electrode (black or colorless in the images), parallel and facing each other in each
cell. The electrodes should be touching the bottom and standing up (2’’ as the height).
The electrodes must be close (about a finger width apart) but not touching each other.
Gently compact the soil to keep the electrodes upright in a stable position. Generally,
the soil will now fill the cells to about half the height. Repeat for three other cells.
4. Measure the voltage of each cell 1 to 4 (Cu – Zn set) in a manner similar to Part I and II
and record in the Data and Calculation section. Connect the leads from the Voltage
Probe to the Cu (+) and Zn (-) electrodes to get a positive potential reading. Make sure
that the electrodes are not significantly disturbed during the measurements or
attachments of the alligator clip leads. Record in Table III.1.
5. Make battery 1-2: Connect cell 1 and 2 in series as shown in Figure 6 and 7. Connect
the negative electrode (Zn) of cell 1 to the positive electrode of cell 2 (Cu) with an
alligator clip lead. Keep the alligator clip lead in the back, away from you so it does not
come in the way as you make the remaining measurements as shown in Figure 7.
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Figure 6 Soil Battery
Figure 7 Soil Battery Lighting an LED
6. Measure the voltage of the battery 1-2 and record in Table III.2 in the data sheet.
Connect the leads from the Voltage Probe to the Cu (+) electrode of Cell 1 with the red
lead and Zn (-) electrode of Cell 2 with the black lead to get a positive potential reading.
7. Make battery 1-2-3: Connect cell 2 and 3 in series (do not disconnect the series
connection for battery 1-2). Connect the negative electrode (Zn) of cell 2 to the positive
electrode (Cu) of cell 3 with an alligator clip lead. You now have a three cell battery 1-23. Keep the alligator clip lead in the back, away from you so it does not come in the way
as you make the remaining measurements.
8. Measure the voltage of the battery 1-2-3 and record in Table III.2 in the data sheet.
Connect the leads from the Voltage Probe to the Cu (+) electrode of Cell 1 with the red
lead and Zn (-) electrode of Cell 3 with the black lead to get a positive potential reading.
9. Make battery 1-2-3-4: Connect cell 3 and 4 in series (do not disconnect the series
connection for battery 1-2-3). Connect the negative electrode (Zn) of cell 3 to the
positive electrode (Cu) of cell 4 with an alligator clip lead. You now have a four cell
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battery 1-2-3-4. Keep the alligator clip lead in the back, away from you so it does not
come in the way as you make the remaining measurements.
10. Measure the voltage of the battery 1-2-3-4 and record in Table III.2 in the data sheet.
Connect the leads from the Voltage Probe to the Cu (+) electrode of Cell 1 with the red
lead and Zn (-) electrode of Cell 4 with the black lead to get a positive potential reading.
Leave the voltage probe connected to battery 1-2-3-4 through the remaining of the
experiment.
You are provided five LED bulbs. Each LED is made of a different material and emits light of
a different color even though they all look clear. LEDs light up if the voltage of the battery
is greater than the forward operating voltage (Vf) listed in Table 2 (Attachment A) and
Figure 4 (Concepts section). Thus, if the voltage of battery 1-2 is 1.6V, it may light up the
red LED but not the yellow or green LED, and so forth. Also the intensity of the light
emitted is dependent on the magnitude of the current passing through the LED. Thus, as
shown in Figure 4, as the voltage applied to a given LED is increased, the magnitude of the
current passing through it increases asymptotically, and thus the intensity of the light will
increase in a similar fashion.
11. Connect the alligator clip leads to one LED bulb (See possible LED colors in Table 4).
Connect a red alligator clip lead to the positive terminal (longer lead on the LED) and a
black alligator clip lead to the negative terminal (shorter lead on the LED).
Table 4 Possible Light Emitting Diode (LED) Appearance Prior to Voltaic Cell Connection
*
Note: You may be given all colorless LED bulbs.
12. Connect the LED to battery 1 (same as Cell 1). Connect the red lead to the Cu(+)
electrode and the black lead to the Zn(-) electrode. Make sure that the electrodes are
not significantly disturbed during the connection of the alligator clip leads. Determine if
the LED lights up and record (Yes/No) it in the Data and Calculations section. If it lights
up, record its brightness (Intensity: D – Dim glow; M – Moderate glow; I – Intense
glow). NOTE: to determine the light on the LED, compare your observation to the
original color/appearance (off mode) of the LED in Table 4.
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13. Connect a red LED to battery 1-2 as shown in Figure 5. Leave the red alligator clip lead
connected to the Cu(+) electrode of cell 1 and connect the black alligator clip lead to
the Zn(-) electrode of cell 2. Determine if the LED lights up and record (Yes/No) it in the
Data and Calculations section. If it lights up, record its brightness (Intensity: D – Dim
glow; M – Moderate glow; I – Intense glow). NOTE: to determine the light on the LED,
compare your observation to the original appearance (off mode) of the LED in Table 4.
14. Repeat step 13 for connection and observation for battery 1-2-3 and battery 1-2-3-4
and record the observations in the Data and Calculations table.
CAUTION: the current flow through a LED is directional and will only flow from the
anode to the cathode, thus if the connection is reversed, the LED will not light.
15. Observe the voltage of battery 1-2-3-4 on the Vernier as you connect, and then
disconnect the LED.
16. Repeat steps 12 thru 15 on all 5 LEDs provided. Record your observations in the Data
and Calculations section.
Note: Use the sample data below as a guide in setting up the soil battery for data collection and
observations. Sample observations below do not represent the actual bulb glow intensities.
EXPERIMENT VIDEO
The instructor will explain this part of the experiment using Figures 6 and 7, and Table 4 above
instead of showing a video.
Part III SAMPLE DATA
Data 1 Sample Cell Potential Reading: BATTERY 1
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Data 2 Sample Cell Potential Reading: BATTERY 2
Data 3 Sample Cell Potential Reading: BATTERY 3
16
Data 4 Sample Cell Potential Reading: BATTERY 4
Data 5 Sample Observation of LED Light Activation and Intensity: BATTERY 1
17
18
19
Data 6 Sample Observation of LED Light Activation and Intensity: BATTERY 1-2
20
21
Data 7 Sample Observation of LED Light Activation and Intensity: BATTERY 1-2-3
22
23
24
Data 8 Sample Observation of LED Light Activation and Intensity: BATTERY 1-2-3-4
25
26
Table Sample Data 2
Sample Cell Potential of Battery 1-2-3-4 as LED is
Connected and Disconnected
LED
No.
Color
1
Red
2
Yellow
Voltage (V)
On Disconnection
On Connection
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3
White
4
Blue
5
Green
LAB SAFETY AND WASTE DISPOSAL
Waste Disposal:
Collect the electrolytes from Part I and II in a beaker and dispose in inorganic waste bottle
labelled “Inorganic Waste” in the Satellite Hazardous Waste Accumulation area. Dispose
of the soil electrolyte in the garbage. Wash, dry and collect all the solid electrodes and
place them in containers as instructed by your instructor.
Lab Safety:
Wear the appropriate Personal Protected Equipment (PPE). Read all Safety Data Sheets
(SDSs) provided by instructor. Pay attention to the safety precautions mentioned in the
procedure and by the instructor. Wash your hands thoroughly with soap or detergent
before leaving the laboratory.
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Bibliography
AspenCore, Inc. (2019). The light emitting diode. Retrieved November 1, 2019, from
https://www.electronics-tutorials.ws/diode/diode_8.html
Randall, J. and Volz, D. (2017). Advanced Chemistry with Vernier, 3rd edition. Electrochemistry:
voltaic cells. Beaverton, OR: Vernier Software and Technology.
Tro, N. (2017). Chemistry: A Molecular Approach. 4th Edition. Boston: Pearson.
Wikipedia (2019). Light-emitting diode. Retrieved November 1, 2019, from
https://en.wikipedia.org/wiki/Light-emitting_diode
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ATTACHMENT A
Table 1 Standard Reduction Electrode Potentials of Some Reactants at 25 oC
30
Table 2 LED Characteristics
Semiconducting
Material
GaAs
GaAsP
GaAsP
GaAsP:N
AlGaP
SiC
GaInN
Wavelength
Color
Vf * @ 20mA
>760 nm
630-660 nm
605-620 nm
585-595 nm
550-570 nm
430-505 nm