MBEYA UNIVERSITY OF SCIENCE AND TECHNOLOGY

COLLEGE OF SCIENCE AND TECHNICAL EDUCATION

DEPARTMENT OF APPLIED SCIENCES

SEMESTER II 2023/2024

BACHELOR OF LABORATORY SCIENCE & TECHNOLOGY

NSLT 07406: CHEMISTRY PRACTICAL II
UQF 8: SECOND YEAR
PRACTICAL HAND-OUT
Practical 5: Synthesis of Aspirin
Section 1: Purpose and Summary
 Conduct a chemical reaction to produce aspirin.
 Separate the aspirin from the reaction by-products using vacuum filtration.
 Analyze the aspirin and estimate its purity.
Acetylsalicylic acid, commonly known as aspirin, is the most widely used drug in the world
today. Its analgesic, antipyretic, and anti-inflammatory properties make it a powerful and
effective drug to relive symptoms of pain, fever, and inflammation. Salicylic acid, whose name
comes from Salix, the willow family of plants, was derived from willow bark extracts.
Hippocrates the ancient Greek physician, as well as Native Americans before Columbus’ time,
prepared willow bark teas as headache remedies and other tonics. In modern times, salicylic
acid is administered in the form of aspirin which is less irritating to the stomach than salicylic
acid.
To prepare aspirin, salicylic acid is reacted with an excess of acetic anhydride. A small amount
of a strong acid is used as a catalyst which speeds up the reaction. In this experiment, sulfuric
acid will be used as the catalyst. The excess acetic anhydride will be quenched (reacted) with
the addition of water. Overall, the reaction takes place between a carboxylic acid and an acid
anhydride to form an ester.

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Aspirin is not very soluble in water so the aspirin product will precipitate when water is added.
Some of the other compounds, acetic anhydride and acetic acid, dissolve in water, but salicylic
acid is only slightly soluble in cold water. Vacuum filtration will separate the crystalline aspirin
away from everything else in the reaction mixture except for any salicylic acid that did not
react.

The aspirin should be analyzed for the presence of any contaminating salicylic acid. In the final
part of today’s lab, you will take a small amount of your aspirin and test it with iron(III)
chloride (FeCl3). FeCl3 reacts with phenols (alcohol groups attached to aromatic rings) to
produce colored complexes. Notice that salicylic acid contains the phenol functional group but
aspirin does not. Therefore, the more salicylic acid that contaminates your aspirin, the darker
the color will be with FeCl3.

Procedure
Part 1: Synthesis of Aspirin
1. Prepare a boiling-water bath by filling a 600-mL beaker with about 400 mL of tap water.
Put the beaker on the hot plate.
2. Weigh out about 2.1 grams of salicylic acid on a piece of weighing paper. To do this,
first tare a piece of weighing paper (This means to put a piece of paper on the balance
and press the zero or tare button. Add some salicylic acid a little at a time onto the
weighing paper until you have about 2.1 gram of it on the paper. It is OK to weigh a
little extra mass (do not return excess salicylic acid to its container as this might
contaminate the entire amount if your spatula is not perfectly clean). Record the mass
of the solid. Place this solid into a 125-mL Erlenmeyer flask (other sizes may be
acceptable).

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3. In the fume hood, measure out 4 mL of acetic anhydride in a small graduated cylinder
and add it to the flask. From this point on, keep your flask under the hood, because it
now contains acetic anhydride (the vapors of acetic anhydride are very irritating).
4. Add about 5 drops of concentrated sulfuric acid. This will be the catalyst for the
reaction.
5. Place the flask in the boiling water bath and clamp it in place.
6. Heat the reaction for at least 15 minutes.
7. Put at least 60 mL of laboratory water into a 150mL beaker (or similar size). Then put
this beaker in an ice-water bath. Use this cold water in steps 9, 12, 13, and 14 below.
8. After the reaction has heated for at least 15 minutes, remove it from the boiling water
bath.
9. In 1 or 2 mL portions, add about 10mL cold water to the reaction, swirl the reaction
between each 1-2mL portions. This water will react and destroy any remaining acetic
anhydride.
10. Put the flask with the reaction into an ice water bath. Crystals of aspirin should
form. Chill for at least 10 minutes in order to obtain the maximum amount of crystalline
product
11. Collect the aspirin crystals by vacuum filtration.
To prepare a vacuum filtration set up:
a) You will need a Buchner funnel, a clean 250-mL vacuum filter flask, a filter adaptor
(often a one-hole stopper), a pre-cut filter paper and Beaker #2. WEIGH the dry filter
paper.
b) Place the filter paper on the Buchner funnel. The filter paper should fit snugly and
cover all the small holes in the funnel.
c) Using the filter adaptor, place the Buchner funnel on top of a vacuum filter
flask. You might need to use a clamp and a ring-stand to keep the setup intact and
upright.
d) Connect a vacuum tubing hose to the vacuum filter flask and to the vacuum supply.
12. Add about 25 mL cold water to the reaction and mix well.
13. Filter the crystals. Use a few milliliters of cold water to transfer as much solid as
possible.
14. Disconnect the vacuum temporarily and add 15 mL of cold water to the crystals in order
to rinse them. After the crystals have soaked in the cold water for a few seconds,

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 reconnect the vacuum to filter once more. Leave the vacuum on for several minutes to
air dry the crystals.
15. With the vacuum on, stir the crystals carefully to help dry them. Do this for about 15
minutes.
16. Turn off the vacuum and weigh just the filter paper with the crystals. If you use a watch
glass to carry the filter paper and crystals to the balance, be sure to tare the weight of
the watch glass (in other words, don’t include the weight of the watch glass with your
filter paper and crystals.)
17. Clean up: Be sure to keep the aspirin for Parts 2 & 3. Throw away the filter paper in
the trash. Dispose of the filtrate (the aqueous solution in the vacuum filtration flask)
down the drain with plenty of water. In the fume hood, rinse with water all pipets and
glassware in contact with acetic anhydride and pour the rinse water down the drain in
the fume hood.
Part 2: Analysis of Aspirin (How pure is it?)

1. Obtain 3 test tubes and about 1 mL of ethanol (CH3CH2OH) to all 3 test tubes. Also
add 1-2 drops of 1% FeCl3 solution to each test tube. Be sure all 3 test tubes have
the same amounts of reagents.
2. Label the first test tube “Aspirin” and add to this test tube a few of your aspirin
crystals that you made today. Stir well.
3. Label the second test tube “Salicylic Acid” and add a few crystals of Salicylic Acid
to this second test tube. Stir well. This is your Positive Control. The positive
control shows you what to look for when a test is positive: in today’s analysis, the
positive control shows you a dark color because it contains unreacted salicylic acid.
4. The third test tube is your Negative Control. Do not add any crystals to this test
tube! The negative control shows you what to look for when a test is negative: in
today’s analysis, the negative control shows you a color when no contaminating
salicylic acid is present.

Compare the colors among the 3 test tubes and describe their similarities and differences.

Part 3: Analysis of Aspirin (How stable is it?)

1. Obtain a dry test tube and add about a pea-sized amount of your aspirin crystals
(borrow some relatively pure crystals from a classmate if your aspirin was highly
contaminated with salicylic acid).

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 2. Tap the crystals to the bottom of the test tube. Use a test tube clamp to hold the test
tube.
3. Ignite a Bunsen burner.
4. Pointing the open end of the test tube away from you, GENTLY heat the bottom of
the test tube until the crystals melt. DON”T burn the aspirin.(Fig.5)
5. Remove the test tube from the flame and carefully WAFT any odors from the open
end of the test tube towards you. Do you smell anything? (To waft the odors means
to “fan” the odors towards you, mixing in plenty of air to dilute the fumes.) If you
don’t smell anything, try “pouring” the vapors out of the test tube by holding the
test tube horizontally.
6. Any odor detected may indicate that aspirin has decomposed with excessive heat.
Aspirin may also decompose slowly upon long-term storage in a humid
environment.
7. After the test tube cools, wash the test tube in the sink with soap and water

Calculations
 Mass of salicylic acid (g)
 Calculate the number of moles of salicylic acid (mol)
(Molar mass = 138 g/mol)
 Volume of acetic anhydride (mL)
 Calculate the number of grams of acetic anhydride (g)
(Density = 1.08 g/mL)
 Calculate the number of moles of acetic anhydride (mol)
(Molar mass = 102 g/mol)
 Identify the Limiting Reagent (Name of compound)
 Theoretical Yield of Aspirin (g)
(Molar mass = 180 g/mol)
 Mass of filter paper and aspirin (g)
 Mass of filter paper (g)
 Calculate the mass of aspirin (g)
 Calculate the Percent Yield of Aspirin (%)
Post Lab Questions:
1. Aspirin is slightly soluble in water: the solubility of aspirin in water is 0.33 grams per
100 mL water at room temperature. In today’s experiment, you rinsed your aspirin with

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 at least 50 mL of water. If the water was at room temperature, how many grams of your
aspirin would have dissolved by rinsing it today?
2. Suggest a reason why you were instructed to rinse the aspirin even though it is known
that this will cause some of the aspirin to dissolve and be lost.
3. Suppose that another student performed today’s experiment, but they forgot to air dry
the aspirin during vacuum filtration. If they weigh the aspirin when it is still damp,
how will this affect the percent yield?
4. Suggest a scientific reason why we did not dry the aspirin in an oven today.
5. A desiccant, or drying agent, is often added to a container with medications to prolong
the shelf life. Cotton is sometimes used as a desiccant which is often found inside a
package of aspirin. Moisture in the air (humidity) can cause aspirin to slowly
decompose through a reaction called hydrolysis. The cotton absorbs moisture and
delays hydrolysis. Suggest an easy and quick way to determine if your aspirin at home
has begun to hydrolyze.
6. Suppose two groups of students performed today’s experiment and obtained different
results. One group obtained aspirin in 88% yield which turned dark in the FeCl3 test.
The other group obtained aspirin in 65% yield which produced no color change in the
FeCl3 test. Explain which group was more successful in lab.

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