Exercise 11
Illinois Central College
CHEMISTRY 130
Page 1
Name:___________________________
Beer's Law: Colorimetry of Copper(II) Solutions
Objectives
In this experiment, we will use Beer's Law to determine the unknown concentrations of Copper(II)
solutions by comparing the amount of light absorbed by the unknowns to the absorbtion of light by a
series of known concentrations.
Copper compounds have been used extensively in the treatment of algae in municipal water supply
impoundments. Consequently, recent indications that copper levels in the sediments of these
impoundments are impeding plant growth have led scientists to more closely monitor Copper levels in
natural waterways.
Background
If white light is passed through a solution containing a colored compound, certain wavelengths of light
are selectively absorbed (taken in). The resultant color observed is due to the transmitted light (light
which passes through). Copper (II) nitrate appears blue to the eye. This is because red light is absorbed
and blue light is transmitted (Table 1). The amount of red light absorbed is directly proportional to
the concentration of the copper (II) ions in the solution as defined by Beer's Law. In this experiment
we will measure the absorbance of several
copper (II) solutions.
Table 1. Correlation between wavelength, color, and complementary color in the visible region.
Wavelength, nm
380-435
435-480
480-490
490-500
500-560
560-580
580-595
595-610
610-750
Color (light absorbed)
violet
blue
green-blue
blue-green
green
yellow-green
yellow
orange
red*
Complementary color(light transmitted)
yellow-green
yellow
orange
red
purple
violet
blue
green-blue
blue-green*
*For Copper (II) nitrate the absorbtion maximum is 630 nm.
Exercise 11
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Beer's Law
In 1852, Beer discovered that the transmittance of light decreases exponentially in proportion to the
concentration of the species absorbing the light. The fundamental law regarding the amount of incoming
light absorbed by a sample is known as Beer's Law. For example, a full bottle of cola placed beside a
bottle containing 1/10 cola and 9/10 water will be drastically different in appearance. This is because
there are more molecules causing coloration in the bottle of straight cola than in the diluted bottle. In
other words, more of the visible light is being absorbed by the straight cola than by the diluted cola.
From this example, it seems reasonable that the amount of light absorbed by a sample, denoted by A,
should be proportional to the amount (or concentration, C) of light absorbing molecules in the sample.
Therefore, Beer's Law can be stated most simply as:
A = k x C
where A is the Absorbance
of light by the sample. The constant, k, depends on the path length through the sample (diameter of the
container), the wavelength of the light used, and the type of absorbing sample.
As shown in Table 1., the color of light that a substance absorbs is the "opposite" of the color the
substance appears; the solution has the color of the light that is not absorbed. As you will see, a
measurement of the absorbance, A, of a sample will allow you to find the concentration of the
light-absorbing sample. In other words, you can quantitatively identify chemicals in solution by the
amount of light they absorb. So, according to Beer's Law, a plot of the absorbances vs the
concentrations of several samples should produce a straight line with a slope, k.
Colorimeter
A colorimeter measures the amount of
light passing through a sample; this
intensity of light is known as the
transmittance.
You will use a Colorimeter (a side view
is shown in Figure 1) to measure the
concentration of each solution. In this
experiment, red light from the LED light
source will pass through the solution
and strike a photocell. A higher
Figure 1.
concentration of the colored solution
absorbs more light (and transmits less) than a solution of lower concentration.
The light sources in the colorimeter are light emitting diodes (LEDs). The LEDs emit a range of
wavelengths with a peak, or most intense, wavelength near the center. The peak wavelengths for the
colorimeter LEDs are 430 nm, 470 nm, 565 nm, 635 nm for the violet, blue, green and red colored
LEDs, respectively. Due to the nature of LEDs, it is incorrect to assume that the light emitted by two
Exercise 11
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LEDs will generate a third color. Therefore, any practical use of the colorimeter will involve only one
LED at a given time.
Since the photocell detector simply changes resistance in proportion to the intensity of the light that
strikes it, we can use the current that passes through the cell to determine the %Transmittance of the
sample where
outgoing light
%T =
incoming light
×100
sample current (microamps)
or
Unfortunately, %T is not linearly proportional
to Concentration. As stated before, it is an
exponential relationship. However,
Absorbance of light by the sample is linear
with concentration. If the current reading (in
microamperes) for the photocell without an
absorbing specimen in the path is Io and the
current reading with an absorbing sample in
the path is I, (Figure 2.) then the absorbance
of the sample is defined as:
I
A = log( o ) or
I
%T =
blank current (microamps)
Clear Cuvette
Light-emitting
Diode (Source)
× 100
Io
CdS Cell
(Detector)
Sample
A = log( 100 )
%T
Connecting the Colorimeter
Connect the Vernier Colorimeter to the
GoLink USB interface and connect the
GoLink to the USB input on your computer.
From the Menu Bar select File/Open and click
on the folder Chemistry with Computers.
Open the file Beer's Law.cmbl. You should
now see the window displayed here.
Right mouse click anywhere on the graph and
choose Graph Options from the pop-up menu. Select the axes
options tab and change the x-axis scaling to 0 for the left to 0.15 for
the right.
Use the arrow buttons on the colorimeter to select the
635 nm LED.
Select a single cuvette to use for both your blank and your samples
for this experiment.
I
Figure 2.
Exercise 11
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Procedure
Preparation of your "Standards"
1. Label five clean test tubes A through E. Fill test tube A approximately 2/3 full with
0.125 M Cu(NO 3)2. Using a 5.00 mL pipette, transfer 5.00 mL of distilled water into test tubes B
through E.
2. Pipette 5.00 mL of solution A into test tube B and mix well. Take care not to lose any of the
solution during mixing. In a similar fashion, pipette 5.00 mL of solution B into test tube C; then
5.00 mL of solution C into test tube D; then 5.00 mL of solution D into test tube E.
3. Calculate the Molarities of each of your standards and record them on the report sheet. Note
that each successive dilution cut the molarity in half.
4. Label three 50 mL beakers "Unknown 1 through 3" and obtain 10 mL samples of the three
unknown Cu+2 solutions.
5. Fill one of the cuvettes with distilled water to serve as a "blank". The blank contains all the
constituents used in the analysis except the substance to be measured. We can assume then that
the difference in the color between the blank and the sample is due only to the substance to be
measured. Distilled water is the reference blank for this experiment.
6. Insert the cuvette containing the distilled water into the opening of the colorimeter. Note that the
cuvette is "ribbed" on two sides. IMPORTANT: Be certain that the light path is passing
through the CLEAR sides of the cuvette facing the arrow at the top of the cuvette slot.
Close the lid of the colorimeter (to keep out stray light) and press the "CAL" button on the
colorimeter to calibrate it. Release the CAL button when the red LED begins to flash. When the
LED stops flashing, the calibration is complete and your unit is ready to collect data.
7. Click
and with the blank still in the colorimeter, click the
button. You will be
prompted to enter a molarity for the sample. Enter 0.0 for the molarity of the blank. Click
OK.
8. Remove the cuvette from the colorimeter and empty it. Fill the cuvette with the copper(II)nitrate
solution from tube E (your most dilute standard.), insert it in the colorimeter, and close the lid.
Allow a few seconds for the Absorbance reading to
stabilize. Click the
button and enter the molarity
for the copper solution in tube E. Continue this same
process until all of the known standards have been
measured, working your way toward the highest Molarity.
9. Once you've finished reading your standard solutions,
from the Analyze Menu, choose Linear Fit. (Or click
on the Linear Fit icon found on the Toolbar.) Your graph
should now look the one displayed here.
Exercise 11
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10. Now fill the cuvette with the first Unknown solution. As soon as the Absorbance reading
stabilizes, choose Analyze from the Menu Bar and select Interpolation Calculator. This
should create a dialogue box on your graph indicating the Molarity of your first unknown. Click
and drag this dialogue box to a vacant area of the graph.
11. Now fill the cuvette with your next unknown and repeat the Analyze/Interpolation Calculator
procedure. Move the dialogue box to another area of the graph.
12. Fill the cuvette with the third unknown and
repeat the Analyze/Interpolation
Calculator procedure. Once all three
unknowns have been analyzed, record the
molarities of your unknowns on your report
sheet. Your graph should now look like the
one shown here.
13. Print a copy of this graph to be attached
to your Report Sheet.
14. Exit LoggerPro.
Exercise 11
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Exercise 11
Illinois Central College
CHEMISTRY 130
Name:________________________
REPORT SHEET
Beer's Law
Standards
Sample
Molarity
%Transmittance
Absorbance
Blank
0.0 M
100
0
A
B
C
D
E
Unknowns
Sample
1
2
3
Absorbance
Molarity
Page 7
Exercise 11 Page 8
Exercise 11 Page 9
Illinois Central College
CHEMISTRY 130
PRELAB: Exp.11
Name:___________________________
Beer's Law
SHOW YOUR WORK
1. A substance that absorbs light at 495 nm appears to have what color? (refer to Table 1.)
2. Referring to the colorimeter in this experiment, if a sample transmits sufficient light to cause a current
of 488 microamperes in the photocell compared to a blank solution that allows a current of 622
microamperes, what is the %Transmittance of the solution?
3. What would the Absorbance value be for the solution in problem #2?
4. Why does the procedure for measuring the concentration of a solution photometrically require
the use of a "blank"?
Exercise 11 Page 10