Aspirin – Part I II & III: Synthesis, Quantitative
& Qualitative Analysis
March 30, 2011
Matt Smith
Bradley LaBoon
Introduction:
Aspirin holds a history that dates back to the days of Hippocrates, where he left medical
records of using powder from the willow bark to ease headaches, pains and fevers. According to
Nick Henderson, Executive Director of the Aspirin Foundation (2009), it was not for years until
the salicin within this willow bark was found to be the ingredient to ease such pain. Many
scientists extracted this salicin over the years until it was titled the name salicylic acid, in its pure
form. It wasn’t until 1900 that aspirin was marketed by Felix Hoffmann who, at the time, worked
for Bayer. At first, aspirin was sold as a powder, where as in 1915, the marketed drug switched to
a tablet (p.12).
This drug has become a valued medical innovation with the help of the salicylic acid
ingredient. Aspirin has been found to treat more than just pain. According to Tash Hughes, of
Word Constructions, the drug has been proven to reduce swelling/inflammation, reduce the
severity of a heat attack, aid recovery after cardiovascular surgery, and even treatment of several
rheumatoid diseases (p.23). Many more uses are tested daily for success, as the salicylic acid
within the drug proves to aid in more than just headache relief. Due to the fact, however, that
aspirin is considered an analgesic, there are possible stomach upsets, and very slight kidney
problems. Despite these side effects, aspirin proves to work wonders in the medical field.
Purpose:
The purpose of this experiment was to first complete an esterification reaction with
salicylic acid and acetic anhydride to produce acetylsalicylic acid in order to observe the purity of
an aspirin sample. Several samples will be produced using varied ingredients in order to examine
the purity of each, noting how each differs from the commercially produced aspirin samples.
From here, several samples of given aspirin will react with an iron (III) chloride solution to
determine the purity of the ‘raw’ and re-crystallized experimental samples produced in the
previous step. The two commercial aspirin samples will also be observed for purity and the results
will be compared to each other.
Method/Procedure:
In Part I of the experiment, salicylic acid and acetic anhydride were reacted to produce
acetylsalicylic acid, or in other words, aspirin:
To begin this, a 6-gram sample of salicylic acid was recorded and combined with 12 mL
of acetic anhydride along with a few drops of sulfuric acid. The reagents were mixed thoroughly
and placed in a hot water bath for a time for 20 minutes. From here, the heated reaction solution
was allowed to cool to room temperature, and 50 mL of distilled water was added to form crystals
before going in an ice bath. The solution then underwent a vacuum filtration in order to separate
the formed crystals from the liquid. The crystals were let to dry and ground up before half were
placed into a vile labeled ‘raw.’ This is when recrystallization occurs in order to improve the
purity of the aspirin sample. The rest of the raw crystals were then dissolved with small amounts
of ethanol and then placed in the hot water bath again. Here, the dissolved crystals run through
the same crystallization procedure as before. However, this process is not rushed, as a good
recrystallization takes time to produce a high quality product. Once finished, these crystals are
placed into a vile labeled ‘pure.’
In Part II, a series of three tests were taken to provide information about the presence of
impurities. The samples synthesized in Part I, were likely to contain small amounts of impurities,
including less in the recrystallized sample. To test other samples, a TLC analysis was taken.
Marks were made along the top and bottom of the TLC plate along with 5 marks on the bottom
for each of the 5 samples (salicylic acid, new commercial aspirin, old aspirin, raw aspirin, &
recrystallized aspirin). A few crystals of salicylic acid were placed on a watch glass along with <1
mL of ethanol. Using a micropipette, a spot was placed on the TLC plate, along the bottom.
This spotting process was repeated for each sample and allowed to dry on the plate. A small
amount of eluent solution was placed in a beaker along with the TLC plate and covered until the
solution reached the end line. The positions of each spot was observed under black light and
recorded. For a second test, five small amounts of each sample were placed in 5 test tubes, each
including 5 mL of ethanol and a few drops of 1% ferric chloride solution. The colors were
observed, with a purple color indicating the presence of salicylic acid. For the final test, melting
point tubes were prepared for each sample and placed in the apparatus. The melting points were
observed and recorded, looking for a range of temperatures to indicate purity.
In Part III, the purity of the aspirin sample was determined through a quantitative
version of the iron (III) chloride test already performed. A series of absorbances of known
concentrations of salicylic acid can be compared to the absorbance of the experimental samples.
To begin, a small sample (.12g - .15g) of salicylic acid was weighed and placed into a flask and
dissolved with 10 mL of ethanol. This was diluted to the 100 mL line with distilled water and
mixed thoroughly, which acted as the stock solution. A series of solutions was prepared using
indicated volumes of stock solution and diluted with a 0.02 M iron (III) chloride solution to a
certain total volume and mixed. To prepare the unknown solutions, 0.15 grams of each sample
(raw, recrystallized, new store-bought aspirin, & old store-bought aspirin) were placed in separate
100 mL volumetric flasks. 10 mL of ethanol was added to each to dissolve the sample and diluted
to the mark with distilled water. 5.00 mL of this solution was place in a 50 mL flask and then
diluted with the iron (III) chloride solution used earlier. These 8 solutions were then placed into
separate cuvettes, about 2/3 full. Using a ‘blank’ of the iron (III) chloride solution, the
spectrometer was set using six parameters. After running with the first solution, the wavelength
and absorbance of the peak on the scan was recorded and entered before running each of the
cuvettes.
Part I:
Mass of salicylic acid
Mass of sample
Raw observations
Pure observations
6.330g
7.195g
One, moist clump of off-white material
Many pure, white crystals that appear dry
in powder form
Part II:
Sample
Range
Raw
Pure
Old Aspirin
New Aspirin
Salicylic Acid
Qualitative Observations
Dark purple
Medium shade purple
Light BurntOrange/Yellow
Light BurntOrange/Yellow
Deep purple
Melting Point
121.0-123.9
129.4-134.0
133.5-134.8
131.0-134.4
158.9-160.2
Data for Preparation of Standard Solutions
Mass of salicylic acid
0.145g
Volume of Stock Solution Prepared 100.00 mL
Standard
Solution
Volume of Stock Solution
Used
(mL)
Final Volume of Standard
Solution
(mL)
Absorbance
A
5.00mL
50.00mL
1.086
B
5.00mL
100.00mL
0.589
C
1.00mL
50.00mL
0.212
D
1.00mL
100.00mL
0.133
Calculations
EQUATION
EXAMPLE
!"## !"#$%&#$% !"#$
!"#$% !"##
Moles salicylic
acid
Molarity of
undiluted
standard
solution
!. !"#$
!
!"#. !"
!"#
= !. !!"!#$%
!"#$%#&'(#&'%'&)
. !""#
!. !!"!#$%&
. !""#
= !. !"# !
! ! !! = ! ! !!
. !"#$ ∗ . !!"
= !!
∗ . !"#
!! = !. !!"#$
Molarity of
diluted standard
solution
Results
Solution
Concentration
A
0.0015M
B
0.00075M
C
0.0003M
D
0.00015M
Graph of absorbance vs. concentration (including trend line):
1.2 1 0.8 y = 717.28x + 0.0208 R² = 0.99714 0.6 Series1 Linear (Series1) 0.4 0.2 0 0 0.0005 0.001 0.0015 0.002 Data for Unknown Solutions
Sample
Sample Mass
Solution Absorbance
Raw Experimental Aspirin
0.153 g
1.190
Recrystallized Experimental Aspirin
0.160 g
0.369
Commercial Aspirin – New
0.150 g
0.024
Commercial Aspirin – Old
0.147 g
0.239
Calculations – using Raw Experimental Aspirin as the example
EQUATION
!−!
!
!. !"#−. !"!#
= !. !!"#$
!"!. !"
!"#$% !" !"#$%&
!"#$%& !" !"#$%&"'
!. !!"#$ ∗. !"" ! !"#$%&"'
= !. ! ∗ !"!!
Molarity of salicylic
acid in solution
!=
Moles of salicylic
acid in solution
!=
Mass of salicylic acid
in unknown sample
!"#$% =
Mass of aspirin present
in unknown sample
Percent purity of aspirin
in unknown sample
EXAMPLE
!"## !"#$%&#$% !"#$
!"#$% !"##
!"## !! !"#$%& − !"## !" !"#$%&#$% !"#$ !"## !" !"#$%& − !"## !" !"#$%&#$% !"#$
!"##$%#"!&'(
∗ !""
!. ! ∗ !"!! ∗ !"#. !"
0.0221g
!
=
!"#
!. !"#$ − !. !""#$ = !. !"!#
!. !"#!−. !""#$
= !". !%
!. !"#$
Results
Sample
Solution
Concentration
Mass of
Salicylic Acid
Mass of
Aspirin
Purity
Raw Exp. Aspirin
0.0016M
0.0221g
0.1310g
85.6%
Recrystallized Exp.
Aspirin
4.85 ∗ 10!!
0.0067g
0.1531g
95.8%
New Aspirin
4.46 ∗ 10!! 6.16 ∗ 10!!
0.1499g
99.9%
Old Aspirin
3.04 ∗ 10!! 0.0042g
0.1428g
97.1%
Discussion:
The purpose of this lab was to observe and analyze the synthesis reaction of
aspirin. To start out the lab, salicylic acid became the main component of the reaction as
it was added with acetic anhydride. This reaction went fairly smoothly, producing decent
samples of both raw and pure aspirin. The ‘raw’ sample, however, seemed to retain more
water than the other sample and was more of a ‘gummy’ texture. Through
recrystallization of the aspirin, the pure sample turned out to be more ‘powder-like’ and
much whiter in color. Once the samples were made, the tests to analyze them began.
First, was the iron (III) chloride test in which each sample was combined with the iron
solution to observe the color produced. The presence of salicylic acid creates a dark
purple color, varying from sample to sample based on the amount contained. The raw
and recrystallized samples both produce a relatively dark purple color, very different from
the burnt orange colors of the store bought aspirins. The raw should have contained
amounts of salicylic acid within the sample, although the recrystallized sample should
have contained a smaller amount. This could have been due to a faulty recrystallization
and/or the contamination of the samples. A vacuum filtration was used to filter out the
crystals from the solution, which could have not completely filtered correctly. If the
process was rushed, and the crystals were not let to dry long enough, the samples may
have retained any bit of water and/or salicylic acid. The TLC plate reading proved to be
very temperamental as our plate was unreadable. This could have been due to the
extended amount of time the plate was left in the eluent solution. This produced a dark
line of color at the top of the plate, proving to be useless. However, another TLC was
provided for qualitative reasons about the 5 different solutions. From here, the melting
point test went very smoothly as the steps to the procedure were minimal. All 5 samples
exhibited melting point ranges that fell well within the standard ranges. In Part III, a
stock solution was prepared to be used for the rest of the experiment as it contained the
iron (III) chloride and salicylic acid. Since the salicylic acid was measured on a non-exact
basis, the amounts of solute differed from group to group-producing slightly different
results. Although the amount massed out fell within the right amount needed to create
the perfect solution. The process of creating both the standard solutions and the
unknown solutions did not seem to have any obvious problems, as the instructions guided
the process along smoothly. Once read in the spectrometer, each cuvette of the standard
solutions proved to have a correct absorbance according to the actual readings given.
When it came to the unknown samples, the absorbance readings were much different and
had a negative value for each of the four samples. This occurrence may have been due to
the over dilution of the samples. Within the procedure, the process calls for a second
dilution after the first, in order to clear up the color of the solution. This extra process
may have washed out the color completely of each, resulting in a lower absorbance value
than the iron (III) chloride ‘blank’ solution.
In terms of qualitative tests, this experiment showed many varying results. In the
iron chloride test, small amounts of ethanol dissolve each of the samples and then
combined with a few drops of 1% ferric chloride. Since the salicylic acid will react with
the iron (III) chloride solution, a distinct purple color was apparent in the salicylic acid
test tube. The new and old aspirin samples went along with this as they produced a light
orange/yellow color, showing the absence of salicylic acid within the solutions. This tells
us that the store bought aspirins are relatively pure, and do not contain large amounts of
salicylic acid. However, the raw and pure samples both produced a dark purple color,
indicating the strong presence of salicylic acid. This should not have come out with such
a dark color in either of the two solutions. While the raw sample would have some trace
amounts of the acid, the pure should have gotten rid of those through recrystallization.
This means the process was not a success. In turn, this means that both the raw and pure
samples have a low percent of purity. In the TLC plate analysis, our plate came up with
a long, dark band across the top. This means the bands had all dissolved and ran to the
top of the plate, providing no understandable results. The time left in the solution must
have been too long for this to happen. When observing the given TLC plate however,
the salicylic acid column contains no band but only a large spot at the top. This suggests
that the salicylic acid was fully dissolved into the solution and was able to run to the end
of the plate. When it came to the new store bought aspirin, there was no spot at the top.
This should have been the most pure out of all the samples, which shows that there was
not trace of salicylic acid within the sample placed. Each spot there after, from left to
right, grows increasingly bigger. This shows the presence of salicylic acid increases as you
go from new to old to recrystallized to raw samples. Finally, the melting point test
showed us that the samples we had produced or were given seemed to be very precise.
Each sample melting under the heat fell within the normal ranges of each sample. This
proves the identity of each sample as it was given. All in all, the qualitative tests had
several difficulties but seemed to show vital information. While the melting point test
proved to be correct, the other two tests were not consistent. This questions the purity of
the raw and pure samples made in Part I.
In terms of the quantitative tests, the results seemed to be consistent in terms of
purity. The standard solutions created with the given values in the table seemed to give
accurate absorbance values. As they were produced on the day of the Part III
experiment, they did not have any dealings with the other samples made in previous
weeks. They simply acted as a gauge to what the absorbances should have looked like in
the end. The unknown solutions, however, produced negative absorbance values. When
observing these cuvettes before placement into the spectrometer, the colors all seemed to
be extremely washed out. After going back to the lab manual and rereading the
procedure, this could have been due to the second diluting stage of the process. The
experiment calls for a second dilution with distilled water. After it was once diluted with
the iron chloride solution, the distilled water would have most definitely washed out any
color previous to this process. When compared to the blank solution, the color would be
much less, creating a negative absorbance value. While this does not provide much
insight into the problems of Part II, it remains consistent with the faulty misreading’s of
the created solutions. After examination, I would say the second week provided the most
insight into the reliability of my samples. One would imagine the quantitative tests to
provide the most information about an experiment, however the qualitative results
provided direct results able to be compared to one another. Seeing the dark colors of the
produced raw and recrystallized samples proved the presence of salicylic acid in a direct
way. In week three, the actual procedure proved to be faulty, as the second dilution
process washed out the colors of all samples. If this had not been true, the third week
would have provided concrete evidence to the reliability of the samples using absorbance.
References:
Henderson, Nick. "What Is Aspirin - 100 Years of Aspirin." Aspirin Foundation. Mar. 2009. Web.
30 Mar. 2011. <http://www.aspirin-foundation.com/what/100.html>.
Hughes, Tash. "Aspirin Use Article." Word Constructions Writing Services. 23 Apr. 2003. Web. 30
Mar. 2011. <http://www.wordconstructions.com/articles/health/aspirin.html>.