Experiment 7
CHEM 131
Experiment 7B: Spectrophotometric Analysis of Aspirin
In this experiment, you will attempt to determine the amount of aspirin in a commercial aspirin
tablet. The technique involves hydrolyzing the aspirin, and then using spectrophotometric
analysis to determine the amount of resulting salicylic acid.
Background
Many commercially available pharmaceuticals are sold in tablet form. These tablets contain not
only the active ingredient(s), but also some inert binder material. The binder helps keep the
tablet intact until it reaches the desired target location (e.g., stomach), and may also serve to
provide sufficient size for easy handling, if the quantity of active ingredient is very small. One
of the goals of this experiment is to determine the percentages of aspirin and binder in a typical
over-the-counter aspirin tablet.
Since the mass of active ingredient (i.e., aspirin) is listed on the label of the container, you could
simply weigh each tablet, and from this calculate the mass percent aspirin in the tablet.
However, a secondary goal of this experiment is to determine whether, in fact, the aspirin tablets
do contain the stated quantity of aspirin, and also what the typical variability is from tablet to
tablet. Your report should include:
•
Best value (Mean +/- Standard Deviation of 3 tablets) for mass of aspirin
•
Mass percent aspirin in each of the 3 tablets
•
Sample calculations and graphs needed to obtain the results
•
Evaluation of results and discussion of possible experimental errors
Also include in your report the answers to the following
QUESTIONS
1. Write a reaction equation for the basic hydrolysis of ethyl propanoate.
2. Write a reaction equation for the basic hydrolysis of methyl salicylate.
3. Write a reaction equation for the acid-base neutralization of propanoic acid.
Experimental Approach
Esters are capable of undergoing basic hydrolysis, also known as saponification (we will better
understand the reason for this name in Experiment 8). Basic hydrolysis involves using OH– ions
to split an ester back up into its constituent alcohol and carboxylic acid. Because aspirin
(acetylsalicylic acid) has a carboxylic acid functional group, the initial reaction of aspirin with
OH– ions will actually be an acid-base neutralization reaction:
CHEM 131
Experiment 7B
1
O
O
OH
O
-
OH–
+
+
O
H2O
O
O
O
Acetylsalicylic Acid
Acetylsalicylate ion
Upon heating, further reaction with OH– results in hydrolysis of the aspirin into salicylate and
acetate ions:
O
O
-
O
O
O
+
O
-
OH–
+
-
O
OH
O
Acetylsalicylate ion
Salicylate ion
Acetate ion
(“alcohol”)
(carboxylate ion)
(ester)
Recall from Experiment 4B that the salicylate ion reacts with Fe3+ ions to produce a purple
complex ion. This complex ion absorbs yellow light at 530 nm. This means that you can
determine the concentration of the purple complex ions, and hence the salicylate they were made
from, using a Spec-20 colorimeter.
O
O
O
-
O
+
Fe3+
-
Fe
3+
O
OH
H
You will first calibrate the Spec-20 using purple complex ions prepared from known solutions of
salicylate (you may work in pairs to prepare the calibration standards). Then you will measure
the absorbance of a purple solution prepared from your hydrolyzed aspirin. From your data, you
CHEM 131
Experiment 7B
2
will be able to calculate the concentration of salicylate in the solution, the moles of salicylate in
your hydrolyzed aspirin sample, and, if you assume that all the salicylate came from hydrolyzed
aspirin, the mass of aspirin in the tablet.
Procedural Hints
A. Hydrolysis of aspirin tablets
After weighing your 3 aspirin tablets, place each into a labeled test tube, and add 5.0mL of 1M
NaOH. Place the test tubes into a boiling water bath, and heat for 15 minutes.
Because some solid material from the binder will remain undissolved, you will need to filter each
solution into a 100mL volumetric flask. Be sure to rinse the filtrate several times with distilled
water to get all the salicylate into the flask.
B. Spectrophotometric Analysis
You can use the same proportions for complexation that you did in Expt 4B (5.00mL sample +
1.00mL complexing agent + 1.00mL pH buffer), but in this case, the complexing agent will be
0.010M FeCl3 (the excess reagent in this experiment). Additionally, it is necessary in this
experiment to use 0.10M HCl in place of KHP buffer, to avoid precipitation of iron hydroxide.
You will need to prepare appropriate concentrations of salicylate solutions to use as your
calibration standards. You should first prepare a 1.00 x 10-2 M stock solution, from which you
will then prepare your standards. You can determine the appropriate concentration range for
your standards by calculating the molarity of complex present in the standard solutions of Expt
4B (then round off to convenient numbers).
Remember that you will need to dilute your hydrolyzed aspirin solutions (from A above), so that
they will be in the linear range of absorbance. You can use an estimate of the aspirin content of
your tablets (~325mg) in order to calculate an appropriate dilution.
Prelab Questions
1. What mass of solid sodium salicylate is needed to prepare 100.0 mL of 1.00 × 10−2 M sodium
salicylate solution?
2. From Expt 4B, the standard solutions ranged in concentration from 5 to 25 ppm Fe3+. What
is the molarity of colored complex formed from 25.0 ppm Fe3+ with excess salicylate? What
molarity of salicylate ions would you need in order to produce this same molarity of colored
complex, if excess Fe3+ were present? (Show that 5.0 x 10-4 M salicylate is a reasonable
concentration for the upper end of your calibration curve.)
3. 0.325 g of an aspirin tablet was hydrolyzed by heating it with 5 mL of 1 M NaOH. The
hydrolyzed aspirin was filtered into a 100mL volumetric flask, which is then filled to the line
with distilled water. If you wanted to dilute this solution into a (second) 100mL volumetric
flask, in order to get a salicylate ion concentration between 1.0 x 10-4 M and 5.0 x 10-4 M,
what size volumetric pipet would you choose? (Choices are 1, 2, 10, 20, 25, 50, and 100 mL)
CHEM 131
Experiment 7B
3