Absorbance Spectrophotometry:
Analysis of FD&C Red Food Dye #40
Note: there is a second document that goes with this one! “2046 - Absorbance
Spectrophotometry - Calibration Curve Procedure.” The second document has
instructions for determining the slope and intercept of a linear data set. You will need to
do this for the calculations in the second week of this experiment.
Introduction
The purpose of this experiment is to determine the concentration of red dye in an
unknown sample, provided by your instructor. Students will familiarize themselves with
use of volumetric pipettes, volumetric flasks and the absorption spectrophotometer. A
review of molarity calculations will be helpful in performing this lab!
Discussion
Absorbance Spectrophotometry is a commonly used laboratory technique for
determining the concentration of substances in solutions. In this experiment we will
prepare several samples of FD&C Red #40, a commercial food dye, and use them to
calibrate the absorbance spectrophotometer. Using the calibrated spectrophotometer, we
will then analyze a sample of Kool-Aid brand drink mix to determine the quantity of Red
#40 in the sample.
Many compounds are colored due to their absorption of visible light. Our eyes are
naturally capable of detecting many different wavelengths or energies of light, and we
perceive these different wavelengths as different colors, as summarized in the table
below:
Color
violet
blue
green
yellow
orange
red
Wavelength
380–450 nm
450–495 nm
495–570 nm
570–590 nm
590–620 nm
620–750 nm
So, for instance, light with wavelengths between 620 and 750 nm are perceived by the
human eye as “red” light. “White” light, such as that produced by the sun or many
artificial light sources, contains all of the visible wavelengths of light combined together;
our eye perceives this as “white” but it is really a blend of all of the colors of the
spectrum (red, orange, yellow, green, blue, indigo and violet).
When white light passes through a colored liquid, some of the wavelengths are absorbed
and some are transmitted. The result is that that the liquid will appear colored to our eye.
If a liquid absorbs all of the light in the violet, blue, green and yellow wavelength ranges,
but allows orange and red light to pass through, then the liquid will appear orange-red
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because its only transmitting orange and red light. All of the other colors are filtered out
by the molecules in the liquid. On the other hand, a substance that absorbs in the
wavelength range 500 -750 nm will appear blue or purple to our eyes, because all of the
red, orange, yellow and green light have been absorbed.
If a sample absorbs no light within the visible spectral range, then it will appear as a
clear, colorless liquid. Water, for instance, absorbs such a small amount of visible light
that our eyes generally cannot perceive the absorbance; unless you have a very large
volume of water, it appears completely clear. (Please note that water is NOT truly clear it DOES absorb visible light to a small degree. That’s why it gets darker as you move to
the bottom of a pool of water, even if the water is clean).
If a sample absorbs ALL of the visible wavelengths then it is opaque; no visible light will
pass through it at all.
Many aqueous solutions of chemicals are transparent (meaning that they transmit light)
but not colorless; they will absorb certain wavelengths and transmit others, resulting in a
colored solution as described above. Aqueous solutions of copper ion, for instance,
appear blue or green because they absorb light in the range from about 450 - 750 nm.
Aqueous solutions of certain iron ions are orange because they absorb in the range from
about 380 - 600 nm.
To illustrate these ideas, consider the figure below. Light is shown on the left as incident
on a sample and then emerging from the sample at the right. In general, Po, the radiant
power of the incident light, will be larger than P, the power of the emerging light,
because the sample will absorb some of the light. This is referred to as the absorbance of
the solution. Not only is Po > P, but the colors (wavelengths) of the emerging light may
be different than the incident light.
sample
Po
P
b
A simple and logical relationship exists between the absorbance of the solution and the
concentration of the solution; namely, that the more sample molecules we place in the
path of the incident light, the greater will be the absorbance of the light. Low
concentrations will transmit more light (and the color will appear lighter or less intense),
and high concentrations will transmit less light (and the color will appear darker or more
intense). This relationship is the basis of Beer’s Law, usually expressed as follows:
A = bc
(equation 1)
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where A is the absorbance,  is a mathematical proportionality constant called the molar
absorptivity, b is the pathlength of the light through the sample (as shown in the figure
above), and c is the concentration of the absorbing species.
Beer’s Law can be used to determine concentration of a compound in solution if we
measure A and know both  and b, or their product b. The law’s most common
application in chemistry is to liquid-phase solutions; the concentration, c, is usually
expressed in molarity (M) and the pathlength, b, is usually expressed in centimeters. The
constant  is specific to each wavelength and units of “M-1cm-1.” If you look at the units
of absorbance in Beer’s Law, you can see that the value of A is unitless (or
dimensionless).
A spectrophotometer is a simple device that measures the amount of light which is
absorbed at a particular wavelength; in terms of the figure above, it compares P (the light
which passes through the sample) to Po (the light which is incident on the sample). In
theory, absorbances can range from zero (a sample which is totally transparent at the
wavelength tested) to infinity (a sample which is completely opaque at the wavelength
tested). In practice, however, absorbances above about 2 are not measurable; above an
absorbance of 2, the sample transmits so little light that it is difficult to measure. (You
will notice as you perform the experiment that if the reading rises above 1.999, the
display will flash, indicating an error – readings above this value are not reliable).
In very simple terms, the spectrophotometer measures how “dark” the color of a solution
is – and Beer’s Law allows us to mathematically relate the “darkness” of the color to the
concentration of the dissolved solute.
Look again at Beer’s law, above – you will notice that it has the form of a linear equation.
The form of a linear equation is shown below, compared to Beer’s law:
y = mx + intercept
A = (b)c + 0
You will notice that a plot of absorbance on the y-axis vs. concentration on the x-axis is
predicted to be a linear plot, with a slope equal to b, and an intercept of zero. In
absorbance experiments, this is referred to as the calibration curve. The first step in
performing an absorbance experiment is to determine the calibration curve, by preparing
samples with known concentration (x-values) and measuring their absorbances
(y-values). The calibration curve is a description or representation of the mathematical
relationship between A and c. Thus, once the calibration curve is complete, you will be
able to interconvert between A and c.
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For example, the graph below is a calibration curve for iodide ion, I1-. Several solutions
of iodide ion were prepared, and their absorbances measured on the spectrophotometer:
Notice that the intercept of the y-intercept of the line is close to zero – this makes sense,
because if the concentration the light-absorbing compound iodide is zero, then the
absorbance should also be zero. In practice, the intercept will be a very small number,
close to zero, but may not be exactly equal to zero. {Don’t be surprised when you do
your calculations if your intercept is not zero!}. Also, the slope of the line has been
calculated and is shown in the bottom left corner of the figure, along with the calculated
intercept.
Now that the calibration curve is known, the concentration of any iodide solution is easily
determined by measuring its absorbance. For instance, if an unknown solution containing
iodide has a measured absorbance of 0.400, then you can see from the graph that the
concentration of iodide must be around 0.005 M; that’s the x-value that corresponds to a
y-value of 0.40. Rather than approximating the concentration, however, we can
precisely calculate it. Knowing that on this linear plot the slope is 79.12874, and the
intercept is 0.0153, the relationship is:
y = mx + b
ABS = 79.12874 [I1-] + 0.0153
If the absorbance is 0.400, the concentration of iodide is calculated to be 0.00486 M.
Note that calculating the molarity gives a much more precise value than simply
estimating it from the graph. This is why the slope and intercept of the line must be
determined. The calculation process outlined in this example is very similar to the
method you will use to determine the concentration of FD&C Red Dye #40 in an
unknown sample.
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The compound which we will analyze in this experiment is known as FD&C Red Dye
#40, commonly referred to as “Red 40” or “Allura Red.” It is found is hundreds (maybe
thousands?) of food and drink products, including soft drinks, candies, processed fruit
products, yogurt, ketchup and barbecue sauces, and much more. (Do an internet search
for “Red Dye 40” or “Allura Red” if you would like to know more). The compound has
a formula of C18H14N2S2O8.
You will receive a sample of Red 40 with a known concentration during the first week of
this experiment, and use it to calibrate the instrument. During the second week, you will
receive a sample of Kool-Aid powder, which contains Red 40. You will prepare a
solution from the powder and determine the concentration of Red 40 in the solution. The
overall goal of the experiment is to calculate % (by mass) of FD&C Red Dye #40 in the
Kool-Aid powdered mix.
NOTE: Your lab manual also has a discussion of absorbance spectrophotometry, in
Appendix IV, under the heading “Light Absorption”; you can use that as a supplement to
this document if you wish.
Procedure
Week 1, Part A: Solution Preparation
 Proper use of volumetric flasks: You will be preparing your dilutions in volumetric
flasks, a type of glassware that is probably new to you. A volumetric flask has an
advantage over a graduated cylinder because it is more precise; a graduated cylinder is
precise to within one tenth of a milliliter (one place past the decimal), but a volumetric
flask is precise to within one one-hundreth of a milliliter (two places past the decimal).
Because we are measuring small quantities in this experiment, it’s best to use very
precise equipment.
Volumetric flasks are very simple to use, but remember a few precautions:
 Be sure to fill the flask exactly to the fill line, and not above! This is the most
common mistake people make when using volumetric flasks. There is a horizontal
marking on the flask, which is the mark for 250.00 mL. If you overfill the flask, there
is no way to correct that error… you have to discard the solution and start over!
 Be certain to use the right stopper or lid for your flask! Some of the flasks have
screw-top lids, and some will require the use of a rubber stopper. If you use a stopper
that is too small, it will get stuck in the neck of the flask! Make sure you have the
right size! It should fit snugly, but not go down into the neck of the flask.
 After filling the flask, mix the solution thoroughly. Simply invert the stoppered
flask a few times to make sure the solution is thoroughly blended.
In this portion of the experiment, you will prepare four solutions, and collect the data
required to complete a calibration curve, as described in the discussion above.
A “stock” solution of FD&C Red Dye #40 will be prepared for you. The concentration
of this solution is expressed as grams of dye/Liter of solution. Record this concentration,
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which you can convert to a molarity (remember that Red 40 has a formula of
C18H14N2S2O8).
Obtain about 125 mL of the stock Red 40 solution in a clean, dry beaker. Please do not
take more than you need! Any excess will be wasted. Using 5 mL, 10 mL and 25 mL
volumetric pipettes, and volumetric flasks, prepare the following dilutions of the stock
solution:




Standard #1: 10.00 mL of stock, diluted to 250.00 mL.
Standard #2: 25.00 mL of stock, diluted to 250.00 mL.
Standard #3: 30.00 mL of stock, diluted to 250.00 mL. (use 25 mL and 5 mL
pipettes).
Standard #4: 50.00 mL of stock, diluted to 250.00 mL. (use 25 mL pipette,
twice).
Calculate the molarity of Red 40 in each of the above solutions. Recall that
M1V1 = M2V2. These four solutions are referred to as “standards”. In chemistry lab, a
“standard” solution is one for which you know the concentration.
Week 1, Part B: Measuring Absorbance of the Standards
 Proper use of cuvettes: Test tubes look perfectly clear to our eye, but the
absorbance spectrophotometer is much more sensitive at detecting light absorbance than
our eyes; glass test tubes do in fact absorb a small but measureable amount of light.
Also, test tubes can vary in thickness (thicker glass will absorb more light, thinner glass
will absorb less light), and sometimes there are small bubbles or impurities in the glass
from its manufacture, which affect the transmission of light and would affect our
readings.
To avoid these problems, absorbance measurements must be made in special containers
called “cuvettes”. They look like test tubes, but are made of an optically pure substance
(quartz) that absorbs very little light. They are also manufactured to higher specifications
than test tubes, so that all cuvettes will absorb a minimum amount of light. Cuvettes are
expensive and must be handled properly to avoid damaging them, and to ensure that your
readings are accurate. Please remember a few important things about using cuvettes:



Please do not get the cuvettes mixed in with the test tubes in your drawer. When
you finish using a cuvette, rinse it with distilled water and place it on the drying rack
provided next to the spectrophotometer.
Cuvettes are very easily scratched! NEVER use paper towels or scrub brushes on a
cuvette! Wipe them with soft tissues only (Kim-Wipes are available for this
purpose).
In the spectrophotometer, the light will pass through the bottom half of the cuvette.
When handling the cuvettes, you should always pick them up by the top half; oils
from your skin will leave residue on the surface – while this residue is not visible to
the naked eye, it will affect the absorbance of the sample! For the same reason you
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

should wipe the outside of the cuvette with a Kim-Wipe before placing it in the
spectrophotometer.
You will notice that the cuvette has a vertical white line at the top; and if you look at
the sample container on the instrument, you will notice that there is a black line at the
front of the opening. When placing the cuvette in the instrument, the white line on
the cuvette should align with the black line on the instrument. This ensures that you
always put the cuvette in with the same orientation, so that your absorbance results
are reproducible.
When you take your measurements, the cuvette should be at least half-full.
 Data Acquisition: It is good practice to rinse the cuvette with some of the
solution that is to be tested. Take a clean cuvette and fill it with a few mL of the first
standard solution. Discard the solution; Red 40 is non-toxic and can go down the sink.
Before using the spectrophotometer, the 0% T and 100% T baselines must be
adjusted. Your instructor will do this for you, but ensure that this procedure has been
completed before proceeding.
After rinsing, fill the cuvette at least halfway with the solution to be tested; place
the sample in the instrument and ensure that the white line on the cuvette matches up with
the black line on the spectrophotometer. Close the lid and record your absorbance.
Absorbance readings should be three places past the decimal; for the analog
spectrophotometer, you will have to estimate the last decimal (as you would with a
graduated cylinder or burette).
The settings on the spectrophotometer are adjusted either by black knobs on the
front of the instrument, or by buttons on the top (depending on which model of
spectrophotometer you are using). Please do NOT adjust any settings; if you accidentally
move the knobs or change settings, please inform your instructor so the
spectrophotometer can be re-adjusted for use!
Any Red 40 waste or excess can be disposed of in the sink.
Part C: Plotting the Calibration curve (between week 1 and week 2).
Note: you should plot your calibration curve and have it ready when you arrive
for week 2 of the lab! For instruction on how to plot and use your calibration curve, see
the other document for this lab, “2046 - Absorbance Spectrophotometry - Calibration
Curve Procedure”. (This can be downloaded from Dr. Rubini’s website, or you can get a
copy from your instructor). Be certain to include the origin (concentration = 0, ABS = 0)
as one of your data points.
Week 2, Part D: Unknown Determination
You will use absorbance spectrophotometry to determine the quantity of Red 40 in a
sample of Kool-Aid drink mix. Make sure that you work with the same
spectrophotometer this week as you did last week! Not all of the instruments will have
the same calibration curve.
Your instructor will furnish you with a sample of Kool-Aid drink mix. Carefully
measure, into 250 mL beaker, 10 - 12 grams of the drink mix. Add about 150 mL of
water - be careful not to add too much! - and stir to dissolve the powder. This may take a
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few minutes to completely dissolve. When the powder is completely dissolved, transfer
the solution to a 250 mL volumetric flask. Rinse the beaker several times with small
portions of water, to ensure that all of the solution has been transferred to the volumetric
flask. Then carefully fill the flask to the fill mark with water. Cap the flask and invert it
several times to make certain the solution is thoroughly mixed.
The solution of Kool-Aid that you have just made will be referred to on the data page as
Solution 1. Solution 1 contains Red 40, but in a concentration that is too high to be
accurately measured by the spectrophotometer – its red color will be too dark!
(Remember from the discussion that absorbance values over 2 are not reliable). So we
will dilute the solution again to bring it down to a range that is measurable. Pipette
50 mL of the Kool-Aid solution into another 250 mL volumetric flask and dilute it with
water to the mark on the flask. As before, invert the flask several times to ensure a
thorough mixing of the dye. This new solution will be referred to on the data page as
Solution 2.
Locate the SAME spectrophotometer that you calibrated last week. (It may not be in
the same spot in the laboratory!). Check with your instructor to assure that the the 0%T
and 100%T points have been set. Rinse a cuvette with a small amount of your second
Kool-Aid mixture, and discard the rinse. Fill the cuvette at least halfway with your mix,
and determine the absorbance of the sample.
Using the calibration curve which you have prepared, determine the concentration of
Red 40 in the Kool-Aid solution.
Any waste or excess Kool-Aid can be disposed of in the sink.
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Absorbance Spectrophotometry:
Data Page
Name _________________________________________________________________
Lab partner _____________________________________________________________
Spectrophotometer number ____________
Week 1: Absorbance of Standards
The stock solution, which was provided, was prepared by dissolving ____________
grams of Red 40 in water to make _________ L of solution.
Standard
1
Concentration of Red 40 (M)
Absorbance
2
3
4
Questions – be sure to show ALL CALCULATIONS on your calculation page!
1. Calculate the molarity of the standard Red 40 solution that was provided. The
molecular formula for Red 40 is C18H14N2S2O8.
_____________________ M
2. Calculate the concentration of Red 40 in the four standards you prepared. Enter them
in the table above.
3. The pathlength of the cuvettes we used in this experiment is 1.00 cm. Use the formula
for Beer’s Law (equation 1 in the discussion) to determine the molar absorptivity of Red
40. (For the purpose of this question, you may assume that the intercept of the calibration
curve is zero). Report your answer for all four standards; they should be very similar,
although they may not be exactly the same.
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Data Page
(continued)
Week 2: Kool-Aid analysis
Mass of powder mix used: ____________________________ grams
Absorbance of Unknown _____________________________
Slope of calibration curve ____________________________
Intercept of calibration curve __________________________
Questions – be sure to show ALL CALCULATIONS on your calculation page!
4. Based on the absorbance of the Kool-Aid, determine the concentration of Red 40 in
Solution 2.
______________________ M
5. How many moles of Red 40 were in your Kool-Aid Solution 2?
_____________________ moles
6. How many grams of Red 40 were in your Kool-Aid Solution 2?
____________________ grams
7. How many grams of Red 40 were in Solution 1?
(This is a very simple question! What is the relationship between the amount of solute in
solution 1 and solution 2?)
____________________ grams
8. Calculate the % of Red 40 (by mass) in the Kool-Aid powder mix.
______________________ %
Remember to submit a copy of your calibration curve with this report.
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Absorbance Spectrophotometry:
Advance Study Assignment
Name ______________________________________________________
1. A solution is prepared by dissolving 0.1896 grams of FD&C Red #40 in water to
produce 2.500 Liters of solution. Given that the molecular formula of Red 40 is
C18H14N2S2O8, determine the molarity of this stock solution.
_________ M
2. 25.00 mL of the dark red stock solution described question #1 is transferred to a
volumetric flask and diluted with water to 100.00 mL.
a. how many moles of Red 40 are in the 25.00 mL sample that was transferred?
_________ moles
b. After the dilution, what is the molarity of the Red 40 in the 100.00 mL of solution?
_________ M
note: the preceding problem (#2) is a dilution calculation. A convenient formula for
calculating concentrations after dilution is M1V1 = M2V2, where M1 and V1 are the
molarity and volume of the concentrated solution, respectively, and M2 and V2 are the
molarity and volume of the dilute solution.
3. a. The solution described in problem #2 is placed in the spectrophotometer. Based on
the calibration curve shown on the next page, what would you predict the absorbance to
be? Note: like the example given in the discussion, the absorbance should be calculated,
not just estimated from the graph. Notice that the slope and intercept of the calibration
curve are provided in the figure on the next page.
Absorbance = _________
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--------------------------------------------------------------------------------------------------------The following calibration curve was produced for Red 40. The slope was calculated to be
2.1033 x 104. The intercept of the calibration curve is 0.0039.
--------------------------------------------------------------------------------------------------------3b. A sample of powered drink mix with a mass of 0.2548 grams is dissolved in water to
make 250.00 mL of solution. A sample of this solution is placed in the
spectrophotometer and its absorbance is measured as 0.3110. Based on the calibration
curve above, calculate the molarity of Red 40 in the unknown.
_________ M
c. How many grams of Red 40 are in the drink mix solution?
_________ g
d. What is the % by mass of Red 40 in the drink mix powder?
_________ %
4. A sample of powdered drink mix with a mass of 11.22 g is dissolved in water to
make 250.00 mL of solution. 50.00 mL of the resulting solution is transferred to a new
container and diluted to a total volume of 250.00 mL. How many grams of the drink mix
powder are in the second solution? (hint: there is a very simple relationship between the
concentrations of these two solutions... this is a basic one-step calculation!)
__________ grams
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