Download Aspirin Synthesis

Aspirin Synthesis
Your team of scientists has been hired by the drug research division of Santa Monica
Pharmaceuticals, a new start-up company formed by some SMC alumni. You have been asked
to evaluate a simple organic synthesis for Aspirin and to determine if the synthesis merits further
funding by Santa Monica Pharmaceuticals. You will have one week to test the synthesis,
evaluate the purity of the product and make your recommendations.
In many ways, the science of organic synthesis is like finding the solution to a puzzle. The goal
of organic synthesis is always to prepare a target molecule, perhaps chosen because of its
intrinsic scientific importance or, perhaps, because of its economic importance. The constraints
imposed upon the synthesis are what makes it challenging to the synthetic chemist.
For example, often a route is chosen because the starting materials and reagents are available
inexpensively or in sufficient quantities. Often working against this constraint is the fact that, for
efficiency, a synthetic route should involve the least number of steps possible, because each
chemical and physical step reduces the overall efficiency. The choice of the most efficient route
is also affected by other constraints. Sometimes (although not in this lab!), efficient synthetic
routes involve highly reactive or toxic species, even if the product has none of those properties.
Such a route might improve efficiency, but it would also increase the complexity of the synthesis
because of the additional safety precautions that would need to be implemented.
Finally, the end use of the product often dictates the acceptable level of purity of the material.
Typically, chemicals destined for use as pharmaceuticals or in the electronics industry are needed
in extremely pure form. Purification strategies can often be as complex as the syntheses
themselves.
In the first part of the lab your team will make aspirin, purify your product and calculate the
percent yield. In the second part of the lab you will perform quality control experiments on your
aspirin, similar to the kinds of tests that are required by the FDA.
O
C
H
OH
O
O
C
CH 3
H
H
H
Figure 1
Acetylsalicylic acid (aspirin), C9H8O4
Aspirin (acetylsalicylic acid) is a versatile drug that is consumed in huge quantities worldwide. It
is a non-steroidal anti-inflammatory drug (NSAID) with a wide range of physiological effects. At
very low doses, aspirin is used to treat and prevent heart attacks and blood clots. At higher doses
it is used as an analgesic to reduce pain and as an antipyretic to reduce fever. At very high doses,
it is an effective anti-inflammatory agent used to treat rheumatic fever, gout and rheumatoid
arthritis. It is also an anticoagulant, it dissolves corns and calluses, and it provokes loss of uric
acid (a toxin) but promotes retention of fluids in the kidneys. It kills bacteria and induces peptic
ulcers. The exact mechanisms of its pharmacological actions are still under study. In many plants,
salicylates can induce flowering.
In the synthesis you will be evaluating, you will start with salicylic acid and make aspirin using
acetic anhydride according to the following reaction:
CH3
O
O
C O C CH3 +
O
HO C
HO
acetic anhydride
salicylic acid
(C7H6O3)
O
HO C
CH3
O
C OH + CH
3 C O
O
acetic acid
acetylsalicylic acid
(aspirin, C9H8O4)
Note that it is the -OH group of salicylic acid that reacts with acetic anhydride to form an esterlike product. The carboxylic acid group of salicylic acid remains unchanged. Phosphoric acid
will be used as a catalyst in this experiment.
Having synthesized aspirin, you will purify it by crystallization and filtration. Although there
will be some loss of product, good experimental technique will minimize the losses. If your
aspirin were going to be used pharmaceutically it would require even further purification.
The main impurity in the crystallized aspirin will be salicylic acid, which will co-precipitate with
the aspirin if the procedure is done too quickly.
The first method you will use to determine the purity is the melting point of your product. We
will use commercial Mel-temp apparatuses to determine the temperature at which your
synthesized crystals melt. If the aspirin is pure, it will melt sharply at the literature value. If it is
impure, it will be lower than the literature value by an amount that is roughly proportional to the
amount of impurity present.
A more quantitative method of determining the purity of your aspirin is to use absorption
spectroscopy (the same method we used in the manganese lab earlier this semester – you may
want to review this material now). In this method you will react a sample of your purified
product with Fe3+ (aq), introduced as Fe(NO3)3. This will form an intensely purple-colored Fe3+salicylate complex with any remaining salicylic acid impurities, but will not complex with the
aspirin. You will use absorption spectroscopy to measure the amount of the complex formed and
determine the amount of impurity of your aspirin.
Experimental
Your team will need to split into two groups. One group should be responsible for synthesizing
the aspirin while the other group prepares the standard solutions and plots the curve you will
use to analyze your product’s purity using absorption spectrometry.
Part I: Preparation of aspirin
Place about 300 mL of water in a large beaker and begin heating this over a Bunsen burner using
a ring stand and wire gauze. You will need the water to be around 75C.
Fill your squirt bottle with deionized water and place it inside a large beaker of ice water to chill.
Now weigh approximately 2.0 grams (record the exact mass) of salicylic acid into a 50 mL
Erlenmeyer flask. Measure out 5.0 mL of acetic anhydride in your graduated cylinder and pour
it into the flask in such a way as to wash any crystals on the walls down into the bottom. Add
five drops of 85% phosphoric acid to serve as a catalyst.
Caution: both the acetic anhydride and the 85% phosphoric acid can cause chemical burns. Be
sure to handle these substances with caution and wash your hands after use.
Once the temperature of the water bath reaches 75C, clamp the flask in place inside the water
bath and heat for at least 15 minutes. Stir occasionally. Maintain the temperature of the water at
75C throughout this process.
At the end of this time cautiously add about 2 mL of deionized water to the flask to decompose
any excess acetic anhydride. Caution: There may be some hot acetic acid vapors evolved as a
result of this decomposition.
When the liquid has stopped giving off vapors, remove the flask from the water bath and add
20 mL’s more of water. Let the flask cool on the bench for about 5 minutes, during which time
you should see crystals of aspirin beginning to form.
After 5 minutes, place the flask in a beaker of ice water to hasten the crystallization and increase
the yield of products. Allow the flask to cool in the ice water for an additional 5 – 10 minutes.
If crystals are slow to form it may help to gently scratch the inside of the flask with a glass
stirring rod.
Collect the aspirin crystals by filtering the cold liquid through a Buchner funnel using suction.
Be sure to weigh the dry filter paper first. Use the chilled water from your squirt bottle to wash
any remaining crystals from your flask into the funnel while filtering.
Rinse the crystals several times using the chilled water from your squirt bottle to help remove
any impurities. Finally, draw air through the filter for several minutes to help hasten drying of
the crystals.
Place the filter paper under a heat lamp to dry. Allow the crystals to dry for at least 30 minutes.
Weigh the product and then heat for an additional 10 minutes. The crystals are dry when
subsequent weighings remain the same (i.e. all the water has been evaporated). The dry crystals
should have a light flaky appearance. Clumpy crystals suggest that they are still wet and need
further drying. It may help speed drying to break up clumps of crystals using your stirring rod.
Part II: Testing the purity
Once the crystals are completely dry you may test the purity using the Mel-Temp apparatus.
Your instructor will demonstrate the proper use of this device.
You will also need to determine the purity of your crystals using absorption spectrometry. For
this you will need to make a set of five or more standard salicylic acid solutions in the range of
10-3 to 10-4 M. You will need to prepare these standards using dilutions in the same way the
standards were prepared in the Manganese lab (refer to the information below, the Manganese
lab on-line handout and the on-line prelab for this experiment for how to prepare these
solutions).
In this case the samples will be purple and so we will measure all the data at a wavelength of
525 nm.
For this experiment we will use 50.00 mL volumetric flasks. Obtain 2 from the stockroom.
Salicylic acid is not soluble in water, but salicylic acid dissolved in ethanol is. Thus, to make
your stock solution you should first calculate how much salicylic acid you will need, weigh this
amount into a 50.00 mL volumetric flask, and then completely dissolve the salicylic acid in
10 mL’s of 95% ethanol solution before adding water to the mark.
To make up the diluted standards, pipette your calculated quantity of stock solution into a second
50.00 mL volumetric flask. Add 10 ml’s of 95% ethanol to be sure the salicylic acid remains in
solution. Then add 10 mL’s of 0.025 M Fe(NO3)3 solution and fill to the mark with deionized
water.
Measure the absorbance of these standards (what should you use as a blank?) and create a Beer’s
Law plot.
When testing the purity of your aspirin product you will need to weigh out approximately 0.10
grams (accurately record the exact mass) of your aspirin crystals into a clean 50.00 mL
volumetric flask. Then add 10 mL’s of 95% ethanol and completely dissolve the solid product to
be sure that everything goes into solution. Next add 10 mL’s of 0.025 M Fe(NO 3)3 solution and
fill to the mark with deionized water. Measure the absorbance of this solution. If the absorbance
is greater than 1.2 absorption units, you will need to repeat this procedure using a smaller mass
of your product.
You can now use your Beer’s law plot to determine the molarity of the salicylic acid impurity in
this solution, and from this the percent mass of salicylic acid impurity in your product. The
percent purity of your aspirin can be determined by subtracting the percentage of impurities from
one hundred.
Lab Report
The report for this experiment is an abbreviated lab report and need only include: Data, Results,
References, Discussion and Conclusions section. No introduction or method is required. The
report must be type written. A short 2-5 page report is acceptable. Be sure to include your
Beer’s law plot and a table detailing how you prepared your standards.
Discuss the yield and purity of your and include a discussion of error analysis.
Finally make your recommendations to the company regarding the practicality of using this
synthesis to make aspirin and any further testing they may wish to do in this regards.