Introduction:
Earlier this semester you
synthesized aspirin from salicylic acid. Unlike previous
reactions where you assumed a 100% yield, this was an equilibrium
reaction which resulted in less than 100% yield. In addition,
you started with one white compound and produced another white
compound. So how do you figure out the purity of your
aspirin?
There are actually several ways
to determine your product's purity: melting point,
chromatography, mass spectrometry, spectrophotometry, and others.
Of course that last one, spectrophotometry, should ring a bell since we
used it previously to determine the thickness of the copper clad on
newer pennies. It turns out that salicylic acid will react with
iron (III) nitrate to produce a complex that absorbs green light, but
aspirin does not. So you can use this to determine the amount of
unreacted salicylic acid that remains in your aspirin, and ultimately
determine its purity.
Theory:
Note: I have added two videos from the Khan Academy which should give you a better understanding of the theory and practice of spectrophotometry. The first video covers the theory of how spectrophotometry works using both an intuitive and algebraic approach. The second video
works through a standard spectrophotometry problem.
Warning: both videos are approximately 13 minutes long. The
first video is 16MB, and the second is 20MB.
When
salicylic acid is dissolved in water, it produces a salicylate dianion,
which reacts with an acidic solution of iron (III)
nitrate, Fe(NO_{3})_{3(aq)},
to produce a highly colored (violet) tetraaquosalicylatroiron (III)
complex:
The
violet color of the complex
results
from the fact that the complex strongly absorbs green light. When
this green is removed from normal white light, we observe violet
(therefore,
green is the compliment of violet). This absorption of green
light
can be used to quantitatively determine the amount of aspirin present
in
the solution. The more green light that is absorbed, the more
violet
the solution, and hence, the more salicylate is present.
If
greenyellow light with a
wavelength
of 530 nanometer is directed into a solution that contains this aspirin
complex, some of the green light will be absorbed:
As you can see, the intensity of the
green light leaving the sample,
I,
is less than the original intensity of the green light,
I_{0}.
There are two ways of expressing this difference. We can talk
about the fraction of light that was transmitted through the
sample, transmittance (T); or we can talk about the
amount of light that was absorbed by the sample, absorbance
(A). As you can see, one is opposite of the other:
transmittance (T)

absorbance (A)

T
= I / I_{0}

A = log
(I_{0}
/ I) = log (1 / T)

The
inverse relationship between
transmittance and absorbance can best be seen in the following figure:
Notice
that the %T can vary from
0 to 100% whereas the absorbance varies from 2.00 to 0.00 absorbance
units. The more light that passes through the sample, the
higher the transmittance and the lower the absorbance.
Conversely, the less light that
passes through the sample, the lower the transmittance and the higher
the
absorbance.
Unfortunately,
a plot of
transmittance versus concentration does not result in a straight
line. However, a plot of absorbance, versus concentration does
provide a straight line:
In a
typical experiment, several
solutions of known concentration of the salicylate complex are
prepared.
Since the concentration of these solutions is known, they are called
standard
solutions. The absorbance of each standard solution is measured
at
the wavelength of maximum absorption (530 nanometer from the spectrum
above)
using a spectrophotometer. A graph of these absorbance values versus
the
concentration of each of the standards should yield a straight line.
This
relationship is known as Beers' Law:
A = a b c
In
this equation, A
is the absorbance of the solution, a is the molar
absorptivity
(a constant for this complex),
b is the path length of cuvette
(in cm), and
c is the molar concentration of the solution
being measured. If the same cuvette is used to measure all of the
solutions, then a and
b are constant.
This means that the absorbance of a solution is directly proportional
to
the concentration of that solution. Therefore, the molar
concentration,
c,
of a solution can be determined by simply measuring the absorbance, A,
of that solution. Although we are actually measuring the
absorbance of the complex, the stoichiometry of the reaction producing
the complex is 1:1. So, if we know the concentration of the complex, we
know the concentration of the aspirin is the same.
O.K.,
lets work through an
example
to see how all of this theory works. Lets assume that you have access to a "STOCK SOLUTION" of salicyclic acid that has a concentration of 1.98 x 10^{3} M. This
"STOCK SOLUTION" is then
diluted
in varying proportions (aliquots) to produce the standard solutions "A",
"B", "C",
"D", "E", and "F" used to create the Beers' Law plot. Solution "A" is produced by diluting 10.0 mL of
the "STOCK SOLUTION" to 50 mL with Fe(NO_{3})_{3}. The
concentration of salicyclic acid in solution "A" can be found using the relationship:
M_{1}V_{1}
=
M_{2}V_{2}
where M_{1}
is
the
molarity of the "STOCK SOLUTION", M_{2} is the molarity
of the solution "A", V_{1} is the volume of the "STOCK SOLUTION",
and V_{2} is the volume of the solution "A":
(10.0 mL) (1.98 x 10^{3}
M) = (50.0 mL) (M_{2})
Therefore,
the concentration of
standard
"A" is 3.95 x 10^{4} M. Now that you know the
concentration
of standard "A", you can use the spectrophotometer to measure it's
absorbance.
In this example, it had an absorbance of 0.348. Likewise, you can
determine the concentration and absorbance for each
of the other standard
solutions:
Solution

mL of Stock

Concentration

Absorbance

"A"

10.0

3.95 x 10^{4}
M

0.348

"B"

8.0

3.16 x 10^{4} M 
0.289

"C"

6.0

2.37 x 10^{4} M 
0.227

"D"

4.0

1.58 x 10^{4} M 
0.161

"E"

2.0

7.91 x 10^{5} M 
0.082

"F"

1.0

3.95 x 10^{5} M 
0.044

Now you have the data you need to create your Beers' Law plot.
However, it would be a good idea to check your data to make sure it is
consistent before you throw away your "Stock Solution". Remember,
the whole idea behind this experiment is that the absorbance of a given
solution will be directly proportional to the concentration of the
aspirin in that solution. If that is the case, then the
Absorbance of a solution divided by the mL of Stock used to create it
should be very nearly constant. For example, if I divide the
measured Absorbance of Solution "A" (0.348) by the mLs of Stock
solution
(10.0 mL), I obtain a value of
approximately 0..035
Absorbance/mL. Likewise, I obtain values of 0.036, 0.038, 0.040,
0.041, and 0.044 for solutions "B", "C", "D", "E", and "F"
respectively.
Since values are all within about 10% of each other, I am confident in
the data I have collected and am ready to create my Beers' Law
plot. Remember that this is sample data that I have create to
make the Beers' Law plot look good. You may notice that the
higher concentration solutions don't show as much Absorbance/mL as the
lower concentration solutions. This can happen if you use a large
sample of aspirin. If this happen, you will have to throw out the
higher concentration result and only used the lower concentration
results.
Once you have determined the
concentration and absorbance for all five standards, you will plot
these points using an 'XY Scatter' plot (Excel).
Your
Beers' Law plot should look like the one below:
Note that most of the points do not fall
directly on the line.
So, we have asked the software to draw the 'best' straight line through
the data. This is the 'Least Squares Fit' or 'Trend line'.
The plot is fairly straight and has a 'goodness' of fit (R^{2})
of 0.9967, where 1.000 is a perfect fit. It also gives us an
equation for the line which we will use to calculate the concentration
of the salicylic acid remaining in your aspirin sample.
Next you will need to process a sample of
the aspirin you
synthesized previously. Lets assume that you used 0.327 g of your aspirin and
processed
it in exactly the same manner as you did the pure salicyclic acid above. Since
we are looking for the amount of salicylic acid, use the molecular
weight of salicylic acid (138.09 g/mol) to calculate the molarity of
"My Aspirin" solution. You
will
end up with 100.00 mL of a 2.37 x 10^{2 }M "My Aspirin"
solution (assuming it is pure salicylic acid).
You then take 5.0 mL of this "My Aspirin" solution and dilute it to
50.0
mL with Fe(NO_{3})_{3}.
The resulting solution has a
concentration of 2.37 x 10^{3} M (again, assuming it is
pure). You then measure its absorbance and obtain a value of
0.079.
When you plotted your standards (five or six depending on whether you decided use the 10 mL aliquot), you
obtained an equation for
the linear regression equation. In our example, that equation was:
Y = 856.09·X +
0.0169
In this equation, 'Y'
is the absorbance,
'X' is the concentration of the solution, '856.09' is the slope of the
line, and '0.0169' is the yintercept. Since we know the
absorbance ('Y'), we can solve for the concentration ('X'):
X = (Y  0.0169)
/ 856.09
X = (0.079  0.0169) / 856.09
X = 7.25 x 10^{5} M
This is the actual concentration of unreacted salicylic acid
remaining in your
aspirin sample. However, we calculated that if your sample was pure
salicylic acid, it should have a
concentration of 2.37 x 10^{3} M.
This means your aspirin sample actually contains:
(7.25 x 10^{5 }M / 2.37 x10^{3} M) x 100
= 3.06% salicylic acid
Therefore, the remainder, 96.94%,
must be pure aspirin!
Procedure:
Hints for using the cuvette and colorimeter:
 A cuvette have two clear sides (the light passes through these), and two ribbed sides, perpendicular to each other.
 Always handle the cuvette using the ribbed sides. You must avoid fingerprints on the clear sides.
 The cuvette must be clean and dry on the outside. Use a ChemWipe for this. DO NOT use a regular paper towel. This will scratch the clear sides.
 After filling with solution, make sure there are no bubbles. You may have to tap it vigorously to remove them.
 Make sure you check that the colorimeter is working properly by
putting in a cuvette of distilled water. It should have an absorbance
of zero. It is is larger than 0.002, let your instructor know so it
can be recalibrated.
 Make sure you put the cuvette in the colorimeter with the ribbed
side facing you. The light beam travels from right to left in the
colorimeter.
Preparation
of Standards for
the
Beers' Law Plot:
In
this section you will produce
five salicylic acid standards of known concentrations.
Spectrophotometric
determination of each standard's absorbance will be recorded and this
data
will be graphically plotted against concentrations to give a standard
curve
(Beers' Law Plot).
 This lab is very time intensive and you must 'multi task' if you
are going to finish. It is important to
study the procedure before coming to lab and not just 'cookbook'
it.
 Obtain
both a 100mL and a 50mL volumetric flask with their corresponding
plastic caps. Clean them by rinsing several times with distilled
water (Note: DO NOT
use soap to clean them.)
 Thoroughly clean your 150mL and 250mL beakers with
soap and water. Rinse them with distilled water and use paper
towels to remove all of the excess water. It is important that they are dry.
 On the back bench you will find a bottle of 0.02 M Fe(NO_{3})_{3} solution. Fill your 250mL beaker with this solution, you may have to refill it before you are finished with the lab.
 You will also find a bottle labeled "PURE SALICYLIC ACID STOCK SOLUTION" on the back bench.
Use the repippetter to obtain 50 mL of this "STOCK SOLUTION" in you
150mL beaker. (Note: be sure to write down the concentration of this solution).
 Make
sure you clean your 10mL graduated pipette by filling it with the
"PURE SALICYLIC ACID STOCK SOLUTION" and then draining it down the sink.
 Using your cleaned 10mL graduated
pipette,
transfer
a 10.0 mL aliquot into your 50 mL volumetric flask and dilute to the 50 mL
mark with 0.02 M Fe(NO_{3})_{3} solution. Be sure to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
Label the flask as "Solution A".
 Rinse your cuvette with
"Solution A"
and then discard. Refill the cuvette with "Solution A" and
measure
its absorbance.
 Using a 10mL graduated
pipette,
transfer
a 8.0 mL aliquot into your 50 mL volumetric flask and dilute to the 50 mL
mark with 0.02 M Fe(NO_{3})_{3} solution. Be sure to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
Label the flask as "Solution B".
 Rinse your cuvette with
"Solution B"
and then discard. Refill the cuvette with "Solution B" and
measure
its absorbance.
 Using a 10mL graduated
pipette,
transfer
a 6.0 mL aliquot into your 50 mL volumetric flask and dilute to the 50 mL
mark with 0.02 M Fe(NO_{3})_{3} solution. Be sure
to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
Label the flask as "Solution C".
 Rinse your cuvette with
"Solution C"
and then discard. Refill the cuvette with "Solution C" and
measure
its absorbance.
 Using a 10mL graduated
pipette,
transfer
a 4.0 mL aliquot into your 50 mL volumetric flask and dilute to the 50 mL
mark with 0.02 M Fe(NO_{3})_{3} solution. Be sure
to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
Label the flask as "Solution D".
 Rinse your cuvette with
"Solution D"
and then discard. Refill the cuvette with "Solution D" and
measure
its absorbance.
 Using a 10mL graduated
pipette,
transfer
a 2.0 mL aliquot into your 50 mL volumetric flask and dilute to the 50 mL
mark with 0.02 M Fe(NO_{3})_{3} solution. Be sure
to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
Label the flask as "Solution E".
 Rinse your cuvette with
"Solution E"
and then discard. Refill the cuvette with "Solution E" and
measure
its absorbance.
 Check
your data to make sure
your absorbance data is decreasing relative to the decreasing
concentration of each solution. For example, the absorbance for
the 4 mL solution should be half of that for the 8 mL solution and the
absorbance for the 2 mL solution should be half of that for the 4 mL
solution, etc. If you find that the 10 mL solution shows
significantly less absorbance that it should, it is possible that it is
too concentrated and has fallen off the linear portion of the Beer's
Law plot.
 Once you are confident with
your data from the pure salicylic acid, you can dump the rest of
the "PURE SALICYLIC ACID STOCK SOLUTION" down the drain.
Preparation a solution of your aspirin:
 Rinse both the 50mL and 100mL volumetric flasks with distilled water to clean them. Remember: DO NOT
use soap to clean them.
 Now you will create a solution from your aspirin and test its purity against the Beers' Law plot you just made.
 Thoroughly clean a 125mL Erlenmeyer flask with
soap and water. Rinse it with distilled water and use paper
towels to remove as much of the excess water as possible.
 Put the Erlenmeyer flask on the balance and use the 'Tare' button to zero it out. Then add approximately 0.3 g (Note: do not
use more than 0.35 g) of
your aspirin and record the mass to the nearest
0.001 g.
 Wash down the inside of the Erlenmeyer
flask with about 30
mLs of distilled water.
 Now turn your hot plate up about half way and heat this
solution
until it has completely dissolved. Be sure you do not let it
boil! If it boils, your aspirin will decompose and you will lose
purity.
 Once all of your aspirin has dissolved, add another 20 mL of distilled water to the Erlenmeyer flask and allow the
solution to cool
until it is comfortable to touch.
 Quantitatively transfer the
solution of your aspirin to your clean 100 mL volumetric flask and
then
dilute with distilled water to the 100.00 mL mark. Be sure to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
Label the flask as "MY ASPIRIN".
 Since this is a new solution,
make sure you clean your graduated pipette by filling it with "MY
ASPIRIN" solution and dumping it down the drain.
 Using the cleaned 10mL graduated
pipette,
transfer
a 5.0 mL aliquot of "MY ASPIRIN" solution into your 50 mL volumetric flask and dilute to the 50 mL
mark with 0.02 M Fe(NO_{3})_{3} solution. Be sure to
thoroughly
mix this solution by inverting the volumetric flask at least ten
times.
 Rinse your cuvette with "MY
ASPIRIN"
and then discard. Refill the cuvette with "MY ASPIRIN" and
measure
its absorbance.
 You have now obtained all of the data necessary to determine the purity of the aspirin you made earlier in the semester.
 Rinse the cuvette and all of
the glassware you used with distilled water and return them to where
you found them.
 Dispose
of your left over
aspirin in the trash. Remove the label from the test tube, clean
and dry it and return it to the instructor's desk. Use a dry
paper towel to remove any aspirin from the
cork and return it to the instructor's desk as well.
Waste
Disposal. All materials
can be washed down the sink with plenty of water.
Calculations:
 Calculate the molarity of
each
of your
standard solutions, "A", "B", "C", "D", and "E".
 Use Excel to produce
your Beers' Law plot. Enter your concentration and absorbances in
two columns and insert a 'Scatter Plot'. Then make sure to add a
'Trend Line'. This
'Trend Line' is the least squares line through your data. You
will also want to set the plot options to show the equation of the line
and the 'R^{2}'on the graph. You can use this equation to
calculate the
concentration of your aspirin sample by using your absorbance value for
'y' and solving for 'x'.
(Updated 11/8/13 by C.R. Snelling)