Synthesis of Aspirin
Isabella Roig
Department of Chemistry, American University, Washington, D.C. 20016
Date of Publication: February 26, 2014
ABSTRACT: The purpose of this experiment was syn-
thesize acetylsalicylic acid from the reaction of salicylic
acid and acetic anhydride. Following the synthesis of
acetylsalicylic acid, commonly known as aspirin, the
crude product was recrystallized and a TLC analysis was
conducted to analysis the purity of the recrystallized and
crude products. The actual yield of 0.546g of acteylsalicylic acid from 2.015g of salicylic acid, which was
found to be the limiting reagent of the reaction. This
experiment produced a 20.79% yield and, correlatively,
a 79.21% error. The Rf values from the TLC analysis of
salicylic acid, the crude product, and the recrystallized
product were 0.818, 0.590 and 0.897, and 0.606 and
0.879 respectively. Because there were two spots shown
in the TLC analysis for both the crude product and the
recrystallized product, it was concluded that the crude
and recrystallized product were comprised of both salicylic acid and acetylsalicylic acid.
INTRODUCTION
Acetylsaliclic acid is commonly known as aspirin, the anti-inflammatory pain reliever. Aspirin
as it is now known was first produced one hundred and ten years ago. It is naturally found in
plants such as the willow tree and myrtle as salicylic acid. Ancient Greeks, Romans, and Egyptians used willow bark and myrtle to reduce pain
and inflammation. Salicylic acid is very acidic
and painful to digest. In 1853, Charles Fredric
Gerhardt was the first add acetyl to the salicylic
acid. In 1897, Felix Hoffman took the acetylsalicylic acid, purified it, and made it more stable
and easier to digest. In 1899, Berlin trademarked
aspirin. When the first World War took place, the
Allies no longer has access to aspirin. After winning the war, the Allies required the Germans to
release the patent for aspirin. Today, aspirin is
available in more than eighty countries. (1)
Aspirin has two main uses in the body. The first
is being an anti-prostaglandin. Aspirin reduces
levels of prostaglandins, which are chemicals released in the body because of pain, fever, and inflammatory. By decreasing the amount of prostaglandins in the body, aspirin reduces fever, pain,
and inflammation. Aspirin also works as a bloodthinning agent. These two main uses of aspirin
are due to effect of aspirin on the enzyme cyclooxygenase, COX. Aspirin inhibits the enzyme
COX, which stops the formation of prostaglandins and therefore stops inflammation. In addition to hindering the formation prostaglandins,
the inhibition of COX also reduces the ability of
platelets to aggregate. For this reason, increased
bleeding is a common side effect of aspirin and it
is not recommended for pregnant women.(2)
Although the decrease in prostaglandins is very
helpful to stop inflammation, prostaglandins are
also very helpful in the body and necessary. Prostaglandins line the stomach and protect the body
from stomach acid. Aspirin depletes the body’s
store of prostaglandins, diminishing the protective lining on the stomach. This can leads to
stomach ulcers. Prolonged and constant use of
aspirin is very dangerous in this way. (2)
Thin Layer Chromatography, TLC, is a technique
used to separate organic compounds. By separating compounds, TLC can be used to determine
the purity of a substance. There are two phases:
mobile and stationary phase. The mobile phase
flows through the stationary phase. Through capillary action, compounds travel with the mobile
phase and separate due to their different polarity
strengths. In this experiment, silica gel plate was
used as a polar stationary phase. The mobile
phase was a 9:1 mixture of ethyl acetate: methylene chloride, making it predominately polar.
Polar stationary phases better absorb polar compounds. Nonpolar compounds will not travel as
far along the silica plate as polar compounds. The
polarity of a compound is determined by its functional groups. Hydroxyl groups are more polars
than esters. Because acetylsalicylic acid has an
ester groups and salicylic acid has a hydroxyl
group, acetylsalicylic acid will travel farther
along the polar silica gel plate. A UV light is
used to see the distances the spots of the compounds traveled on the plate. TLC results are expressed in Rf values. Rf is the distance the solute
traveled divided by the distance the solvent
front.(3)
In this experiment, acetyl anhydride and salicylic
acid are reacted together while in the presence of
phosphoric acid to synthesize acetylsalicylic acid.
The mechanism of reaction is displayed in Figure
1 below. Following the synthesis of aspirin, it
will be recrystallization as a method of purification. To further test the purification of the recrystallization, a TLC analysis will be conducted to
examine the crude product, the recrystallized
product, and salicylic acid.
Figure 1: Reaction of Aspirin
MATERIALS AND METHODS
Lab coat
Goggles
Gloves
50 mL Erlenmeyer Flask
125 mL Erlenmeyer Flask
400 mL Beaker
250 mL Hot Water Bath
250 mL Ice Water Bath
Acetic Anhydride
Salicylic Acid
85% Phosphoric Acid
Distilled Water
10 mL Ethanol
18 mL Ethyl Acetate
2 mL Methylene Chloride
2 Watch Glasses
Stirring Rod
Vacuum Filter Appartatus
1 110-mm Filter Paper
2 pieces of round Filter Paper
Silica Gel Coated TLC Plate
Spotting Pipettes
UV Light
Part 1: Synthesize Aspirin
2.0g of salicylic acid is measured into a 50mL
Erlenmeyer flask. 5.0mL of acetic anhydride and
5 drops of 85% phosphoric acid solution were
added to the flask. The flask is swirled to rinse
any bits of solid that may be on the inner walls of
the flask. The 50mL flask is partially submerged
in a 70-80°C hot water bath. The mixture is heated for 15 minutes, or until the mixture stopped
releasing vapors. The mixture is also occasionally
stirred while heating. After 10 minutes in the water bath, 2mL of distilled water is added to the
mixture. After 15 minutes of heating in the hot
water bath, the mixture is removed from the bath
and 20mL of water is added to the flask. The
mixture is then cool to near room temperature
and transferred to an ice bath for five minutes.
Crystals of apririn formed while the mixture
cooled. Filter paper was weighed and the mass
was recorded to the nerest 0.001g. Vacuum filteration was set up and the mixture was suctioned
with the vacuum. When the most of the liquid
had been drawn through the funnel, the suction
was turned off and the crystals were washed with
5mL of cold, distilled water. After 15 seconds,
the suction was turned back on and the crystals
were rinsed twice more with 5mL of cold, distilled water. The filter paper and product were
transferred to a tared watch glass. The mass of
the product on the filter paper was recorded to the
nearest 0.001g.
Part 2: Purification of Aspirin
0.5g of the crude crystals were set aside for TLC
analysis. The remaining crude aspirin was added
to a 125mL Erlenmeyer flask. 10mL of hot solvent (ethanol/water) was added to the crude aspirin in a warm water bath until all the crystals
were dissolved. The mixture was not boiled. The
flask was covered with a watch glass, removed
from the warm water bath, and left to cool slowly. When the solution reached room temperature,
it was placed in an ice bath to complete cyrstalization. The crystals were vacuum filtered after 10
minutes in the ice bath. The crystals were rinsed
with two 3mL portions of ice cold deionized water and one 2mL portion of ice cold ethanol. The
crystals were place on a tared watch glass and the
mass was recorded.
Part 3: TLC Analysis
A 400-mL beaker and a watch glass were used as
the developing chamber. A piece of 110-mm filter paper was placed at the bottom of the chamber
to saturated the chamber with solvent vapors.
20mL of a 9:1 ethyl acetate and methylene chloride mixture was added to the 400mL beaker. In
three separate beakers, 3mg of salicylic acid, 3mg
of the crude product, and 3mg of the recrystallized product were dissolved in 5-6 drops of the
TLC solvent. A horizontal line was lightly
marked using a pencil on the silica gel coated
TLC plate approximately ½ inch from the bottom. Three even spaced hashes were marked to
be the origin points for the spotted compounds. A
different spotting pipette was used to spot the
TLC plate with the salicylic acid, the crude product, and the recrystallized product on the three
hashes. The pipette lightly touched the TLC plate
at the hash mark to transfer some of the liquid to
the plate. The spots were no larger than about 1/8
of an inch. The plate was placed in the developing chamber and the watch glass was place on top
of the beaker. The plate was allowed to develop
until the solvent from was approximately ½ inch
from the top. The plate was removed and the solvent front was immediately marked. The plate
was left to dry before it was examined under the
UV light. The spots were circled as they were
seen and their centers were marked. The distance
the solvent traveled and the distance each spot
traveled were recorded.
RESULTS
Table 1: Synthesis of Aspirin
Mass of salicylic acid (g)
Volume of acetic anhydride (mL)
Mass of acetic anhydride used (g)
Mass of filter paper (g)
Mass of aspirin and filter paper (g)
Mass of aspirin synthesized (g)
2.015
5
5.4
0.132
1.35
0.546
This table shows the data that was collected in
this lab. The mass of the aspirin and the filter paper is not equal to the mass of synthesized aspirin
plus the mass of the filter paper. This indicates
that some aspirin was lost, and not a full yield
was collected during the experiment. This is supported by the large percent error calculated below.
Table 2: TLC Analysis
Compound
Distance
Solvent
Traveled
(cm)
Distance Spot
Traveled (cm)
Rf
Value
Crude Product
3.9
2.3 & 3.5
0.589
&
0.897
Salicylic Acid
3.3
2.7
0.818
Recrystalized
Product
3.3
2.0 & 2.9
0.606
&
0.878
The crude and recurstalized products both had
two spots on the TLC Plate. This indicates that
neither were pure and there was more than one
substance in each
Calculations:
Equation 1: Determining the Limiting Reagent:
Molecular Weight of Salicylic Acid:138.12 g/mol
(2.015g Salicylic Acid) x (1 mol/ 138.12g)
=0.01458 mol
Molecular Weight of Acetic Anhydride:
102.09g/mol
(5.4g Acetic Anhydride) x (1 mol/ 102.09g)
=0.05289 mol
Salicylic Acid is the limiting reagent.
Equation 2: Theoretical Yield:
(0.01458 mol Salicylic Acid) x (1 mol Acetylsalicylic Acid/ 1 mol Salicylic Acid) x (180.16g/ 1
mol Acetylsalicylic Acid)= 2.6267g Acetylsalicylic Acid
Equation 3: Percent Yield:
Perecent Yield =(Actual Yield/ Theoretical
Yield) x 100
(0.546g/ 2.6267g) x 100 = 20.79%
Equation 4: Percent Error:
Percent Error: ((Actual Yield- Theoretical Yield)/
Theoretical Yield) x 100
(( 0.546-2.6267)/ 2.6267) x 100= 79.21%
Equation 5: Rf Value:
Rf Value = (Distance Solute Traveled/ Distance
Solvent Traveled)
Rf (Crude Product, Spot 1)= (2.3cm/ 3.9cm)=
0.590
Rf (Crude Product, Spot 2)= (3.5cm/ 3.9cm)=
0.897
Rf (Salicylic Acid)= (2.7cm/ 3.3cm)= 0.818
Rf (Recrystallized Product, Spot 1)= (2.0cm/
3.3cm)= 0.606
Rf (Recrystallized Product, Spot 2)= (2.9cm/
3.3cm)= 0.879
Table 1 displays the raw data from this experiment. This experiment began with 2.015g of salicylic acid and produced a yield of 0.546g of acetylsalicylic acid. Using Equation 1, it was determined that salicylic acid was the limiting reagent.
Equation 3 determined that the theoretical yield
of 2.6267g of acetylsalicylic acid. Equation 4
calculated the percent yield was 20.79%. The
percent error was 79.21%. Table 2 presents the
data of the TLC analysis part of the experiment.
The Rf values of crude product, salicylic acid,
and recrystallized product were 0.590 and 0.897,
0.818, and 0.606 and 0.879 respectively.
DISCUSSION
Esterification is the reaction of an alcohol or
phenol and a carboxyl group in the presence of an
acidic catalyst to form a carboxylate ester. The
particular esterification process in this experiment of forming acetylsalicylic acid from salicylic acid and acetyl anhydride is also an equilibrium process. (4) Le Chatelier’s principle, which is
also known as the equilibrium law, can be used to
predict the effect of a change in conditions on an
equilibrium process. Temperature, concentration,
volume, and pressure are all changes in condition that would cause the equilibrium of the reaction to shift. As the reaction proceeds, the reactants are depleting while the concentration of the
products is increasing. It is also stated in the
principle that if more of the product were added
to the reaction, this would cause the equilibrium
to be shift and the creation of more of the reactants. Le Chatelier’s principle can be used to favor the creation of the product of the reaction,
aspirin, due to the equilibrium shift caused by the
excess acetic anhydride. The principle can also
be used to favor the reactants if the necessary
conditions were adjusted to convert aspirin to salicylic acid and acetic anhydride. (5)
Salicylic acid contains two acidic functional
groups: a phenol group and a carboxylic acid.
The phenol group of the molecule is very acidic
and causing harsh stomach irritation, which
therefore makes it difficult to digest. To ease
stomach digestion of aspirin, the strength of the
acid is reduced by the formation of an ester from
the phenol and carboxylic acid.(6)
Figure 2: Mechanism of Reaction for the Synthesis of Aspirin
purify the crystals. The addition of water separates the crystals from the very water soluble acetic acid. recrystallize the product. The water was
added after heating the mixture because the aspirin would not have formed if the acetic anhydride
had already reacted with water. Water also aided
in nucleophilic substitution at this point in the
recrystallization. The following figure details the
reaction of water and acetic anhydride.
Figure 2 details the reaction occurring between
salicylic acid, phosphoric acid, and acetic anhydride. When the phosphoric acid, acting as a catalyst, attacks the carbon oxygen bond (C=O) of
the acetic anhydride, it gives it a positive charge.
This results in the acetic anhydride being more
prone to nucleophilic attacks. Salicylic acid is the
nucleophile of this reaction. The partial positive
carbon on the acetic anhydride transferred electrons to the oxygen on the phenol group of salicylic acid after it was attacked. This results in the
formation of a tetrahedral intermediate. The –OH
group, which has an attached electrophilic carbon, forms a double bond between the carbon and
oxygen (C=O) by pronating the hydrogen and
donating an electron. This loss of a proton reforms phosphoric acid and produces acetylsalicylic acid.
Because the synthesis of aspirin is a slow reaction, a catalyst is needed. Phosphoric acid acts as
a catalyst to speed up the reaction between Salicylic acid and acetic anhydride. Phosphoric acid
is present in the beginning and end of the reaction. Due to its water solubility, it is easily removed by washing the crystals with cold distilled
water. The phosphoric acid was used to aid the
transfers of hydrogens and electrons. At the end
of the reaction, it was recreated in its original
state.
The reaction was heat at 70-80 degrees Celsius to
speed up the reaction of the synthesis. The heating caused the salicylic acid to dissolve more
quickly as well as increase in solubility. This is a
endothermic synthetisis reaction. Therefore the
increase in temperature causes the reaction ot
move forwrad, favoring the formation of products.. (6)After heating the mixture for 10-15
minutes, water was added to the crude product to
Figure 3: The Reaction Mechanism of Water and
Byproducts
The crude aspirin is dissolved in a mixture of hot
water and ethanol. Water has a polarity of 9. Ethanol has a polarity of 5.2. (7) Of the two solvents,
ethanol is the less polar one. Aspirin has a solubility of 10mg/mL in water and a solubility of
50mg/mL in ethanol. Like dissolves like, therefore less polar solvents dissolve less polar solutes. Therefore water and ethanol are great solvents to help purify aspirin. The crude aspirin
may not have been clean from the earlier recrystallization due to its contact and exposed with the
environment.
Pure aspirin was not made in this experiment.
The Rf values of crude product, salicylic acid,
and recrystallized product were 0.590 and 0.897,
0.818, and 0.606 and 0.879 respectively. Because
the crude and recrystallized product showed two
spots on the TLC plate, it can be concluded that
the crude and recrystallized product were composed of both acetylsalicylic acid and salicylic
acid. The spots of the crude and recrystallized
product TLC analysis correlate to be either salicylic acid or acetylsalicylic acid.
Because the silica was a polar stationary phase, it
better absorbs polar compounds. Nonpolar compounds are not as absorbed as polar compounds
and therefore more freely along the solvent front,
which is polar. Salicylic acid has a hydroxyl
functional group. Acetylsalicylic acid has an ester
functional group and has a larger mass than salicylic acid. Esthers are not as polar as hydroxyl
groups are more polar than esters. Therefore, salicylic acid was more to the silica than the acetylsalicylic acid. This is reflected in the Rf value
of 0.818. The acetylsalicylic acid traveled further
along the TLC plate than the salicylic acid. From
this conclusion, it can be determine that the spot
with Rf values of 0.897 in the crude product and
0.879 in the recrystallized product were acetylsalicylic acid.
The experiment yielded 0.546g of synthesized
aspirin. The theoretical yield for this experiment
was 2.6267g. As seen in the results section by use
of Equations 3 and 4, the percent yield was
20.79% and the percent error was 79.21%. Because the synthesis of aspirin is an equilibrium
process, t is difficult to obtain 100% yield of because a large concentration of the products can
cause the reaction to reverse and create the reactants. In addition to this reason, a few errors
could have contributed to the low percent yield of
this experiment. The percent yield is relatively
low because of possible sources of error. Prolonged exposure to the environment and atmospheric air that may have caused air drying. Product could have also been lost while using vacuum
filtration. Product may have slipped between the
filter top apparatus and the filter paper. Inadequate heating time or high enough temperature
may have not allowed for a complete dissolution,
resulting in a smaller product yield.
CONCLUSION
This experiment yielded 0.546g of acetylsalicylic
acid. The theoretical yield of this experiment was
2.6267g, given that 2.015g of the limiting reagent, salicylic acid, was used. Respectively, the
percent yield and percent was 20.79% and
79.21%. As noted by the small percent yield,
much of the product was lost or not successfully
synthesized to aspirin. The Rf values of crude
product, salicylic acid, and recrystallized product
were 0.590 and 0.897, 0.818, and 0.606 and
0.879 respectively. The solvent used as a mobile
phase for the TLC analysis was a mixture of ethyl
acete and methylene chloride, making it predominately polar. Because the salicylic acid is polar,
along with the mobile phase and stationary silica
gel phase, it had a lower Rf value than the acetylsalicylic acid as it was absorbed more to the
plate. Salicylic acid displayed an Rf value of
0.818, while the two higher spots of the crude
and recrystlization product that are presumed to
be acteysalicylic acid displayed respective Rf
values of 0.897 and 0.879. To prevent further error, the recrystallized product needs to be better
purified. To do so the experiment should be more
closely and entirely conducted under the fume
hood to prevent and decrease exposure to atmospheric air. Also a more temperature controlled
hot water bath could also be helpful in better recrystallization, as well as a bit more distilled water to aid in the dissolution of acetic anhydride
from the crystals.
REFERENCES
1. History of Aspirin; ASPREE: Mansh University, Department of Epidemiology and
Preventive Medicine, 2010.
http://www.aspree.org/AUS/aspreecontent/aspirin/history-aspirin.aspx (accessed Feb 22, 2014)
2. How Aspirin Works; ASPREE: Mansh
University, Department of Epidemiology
and Preventive Medicine, 2010.
http://www.aspree.org/AUS/aspreecontent/aspirin/how-aspirin-works.aspx
(accessed Feb 22, 2014)
3. Williamson, K.; Masters, K. Macroscale
and Microscale Organic Experiments, 6th
ed.; Cengage Learning: Belmont, 2011
4. Clark, Jim. Esterification; ChemGuide,
2003.
http://www.chemguide.co.uk/organicprop
s/alcohols/esterification.html (accessed
Feb 21, 2014)
5. Le Chatelier’s Principle; Purdue University, Department of Chemistry.
http://chemed.chem.purdue.edu/genchem/
topicreview/bp/ch16/lechat.html (accessed Feb 21, 2014)
6. Clark, J. The Effect of Temperature on
Reaction Rate. Chemguide, Oct. 2013,
n.p.
http://www.chemguide.co.uk/physical/bas
icrates/temperature.html (accessed Feb
21, 2014).
7. Bayers, John A. Polarity Index; Chemical
Ecology, 2003. http://www.chemicalecology.net/java/solvents.htm (accessed
Feb 21, 2014).
ACKNOWLEDGMENTS
The author would like to acknowledge the Americam
University Chemistry Department as well as Dr.
Dehghan and TA, Karlena Brown and the students
who participated in the experiment. Acknowlegement
is also extended to those mentioned in the References
section of this paper.
This assignment is my own work and I have cited
all material used in its preparation. This assignment has not previously been submitted at any
other time or any other course. I have not copied
in part or whole the work of other students or
person.
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DATE__________________________________
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