Synthesis and Analysis of Aspirin and Copper (II) Aspirinate by Breckon Pav
General Chemistry II Dr. Soli May 1, 2003 Synthesis and Analysis of Aspirin and
Copper (II) Aspirinate Abstract This project included several experiments,
including aspirin and copper (II) aspirinate synthesis, melting point analysis,
infrared (IR) and ultraviolet (UV) spectra analysis, pH titration, and
conductimetric titration. Each of the five latter experiments was performed to
prove that aspirin and copper (II) aspirinate were indeed synthesized. The
results from these procedures supported this conclusion. Introduction This
multi-faceted experiment involved synthesis and analysis of aspirin
(acetylsalicylic acid) and copper (II) aspirinate. In the synthesis portion of
the experiment, the effects of several catalysts on the speed of an exothermic
reaction were ranked accordingly. The physical properties of aspirin were
analyzed in a melting point procedure, and compared to literature values. A pH
titration was performed to examine the acidic properties of acetylsalicylic
acid. The equivalence point was determined to find the Ka for aspirin. A
conductimetric titration was performed to determine the synthesized compound's
molecular weight. Finally, IR and UV spectra showed the light absorbance
properties of the synthesized product. Theory The primary chemicals analyzed in
this experiment were salicylic acid (Figure 1), acetylsalicylic acid (Figure 2),
and copper (II) aspirinate (Figure 3). Both salicylic acid and acetylsalicylic
acid have the appearance of a white crystal, while copper (II) aspirinate is a
blue crystal. The reaction between salicylic acid (C6H4(OH)COOH) and acetic
anhydride (C4H6O3) yields acetic acid (CH3COOH) and acetylsalicylic acid
(C6H4(CH3CO2)COOH) in an exothermic reaction: C6H4(OH)COOH + C4H6O3 à CH3COOH +
C6H4(CH3CO2)COOH + heat (1)1 The effects of several catalysts (anhydrous sodium
acetate, pyridine, and sulfuric acid) on this reaction were observed. Since the
reaction in question is exothermic, its rate can be roughly estimated by an
increase in temperature. Therefore, the speed of the reaction in the presence of
the different catalysts could be measured based on a 4º C temperature change.
The catalysts can then be ranked according to the magnitude of their influence
on the reaction rate. The reaction between aqueous copper and dissolved aspirin
yields copper (II) aspirinate, a fine blue crystal that precipitates from the
two liquids according to the following chemical reaction: 2 Cu2+(aq) + 4 Asp-H(alcohol)
à Cu2Asp4 (s) + 4 H+(aq) (2)1 IR spectroscopy was performed on copper (II)
aspirinate as well as on the aspirin itself. In addition, an ultraviolet-visible
spectrum was taken of copper (II) aspirinate and compared to that of copper (II)
acetate. If a significant difference between the known compounds and those
synthesized was observed, it was accurately concluded that aspirin and copper
(II) aspirinate were indeed synthesized. Infrared spectroscopy takes advantage
of the vibrations of different groups of atoms at various wavelengths of
infrared light. The spectrometer sends a beam of light through the molecules
that gradually increases in wavelength, affecting different bonds and atoms in
the molecule along the way. The graphic representation of this is displayed as
percent transmittance vs. wavenumber (1/wavelength or cm-1). The graphs of Nujol
(medium) and salicylic acid can be combined to form the graph for
acetylsalicylic acid. Therefore, while the aspirin IR spectrum appears only 50%
similar to each of the others when superimposed, it represents the sum of both.
Ultraviolet-visible spectroscopy (uv-vis) differs slightly with IR spectroscopy
in that it is not generally used to determine a molecule's structure or chemical
composition. Also, the graph plots the absorbance vs. wavelength, where
absorbance is defined as A = -log (%T/100), and %T is the percent transmittance.
The graphs of copper (II) aspirinate and copper (II) acetate should differ
slightly to signify that copper (II) aspirinate crystals were synthesized.
Normally, aspirin melts between 138º and 140º C. The melting point range for
salicylic acid is 158º-160º C. At these temperatures, the crystals' lattice
energy is broken, and they undergo a phase change from solid state to liquid
state. These temperatures are easily attainable and measurable, and therefore
can be attained experimentally. Samples can be contaminated or wet, therefore
altering the experimental temperature ranges. In any pH titration, the endpoint
is defined as the volume of base needed to completely neutralize the acid
present, or vice versa. As base is added to an acidic solution, the H+ ions bond
with OH- ions to form neutral H2O. Na+ + OH- + Asp- + H+ à Na+ + Asp- + H2O (3)
When all of the acid is thus "neutralized," the pH will rise steeply
and the solution will be come increasingly basic. The acid's Ka is defined as Ka
= [H+][A-]/[HA]. A simple way to calculate this value is to evaluate the pH at
the halfway point of the titration. When half of the acid has been neutralized,
[A-] = [HA], so [A-]/[HA] = 1. Therefore, at the halfway point, Ka = [H+], so
pKa = pH2. By performing some short calculations, the value of Ka can be found,
and compared to the literature value. The purpose of the conductimetric
titration is to determine an acid's formula weight. Aspirin, a weak acid, has a
low conductance, while Na+ and OH- both have a high conductance. When the
equivalence point is reached, the concentration of conducting ions will increase
dramatically, and so will the solution's conductance. By determining the number
of moles of base (and acid) present at the equivalence point, the formula weight
can be calculated using the molarity of the standard base used in the titration.
Experimental Several procedures were performed in this experiment, beginning
with the synthesis of aspirin (acetylsalicylic acid). The reaction between
acetic anhydride and salicylic acid was catalyzed in three different ways to
yield acetic acid and acetylsalicylic acid. Approximately 1 gram of salicylic
acid was placed into each of three test tubes for catalytic analysis. To each
tube was then added 2 mL of acetic anhydride. The first catalyst used (tube #1)
was anhydrous sodium acetate (0.2 grams). The compound was stirred, and the time
for a 4º C temperature increase was recorded. To the second test tube was added
five drops of pyridine, and the time needed for the temperature change was also
recorded. Finally, five drops of concentrated H2SO4 were added to test tube #3,
and the time taken for the temperature change was noted. To complete any
unfinished reactions and dissolve any remaining salicylic acid, the three tubes
were placed into a hot water bath for approximately 15 minutes. To allow the
crystals to begin forming, the contents of each test tube were combined in a
125-mL Erlenmeyer flask with 35 mL of deionized water. The flask was swirled
around to help with the hydrolysis of any remaining acetic anhydride, then was
placed in an ice bath. After crystals had formed, they were filtered and
collected with a Büchner funnel and aspirator flask. The Erlenmeyer flask was
rinsed several times with D.I. water to collect any remaining aspirin residue.
Finally, the product was placed in an oven for about 15 minutes to evaporate the
crystals. For several experiments (IR spectra, pH titration, and
electrodeposition), it was necessary to synthesize copper (II) aspirinate
crystals. First, 0.33 grams of copper (II) acetate were dissolved in 50 mL of
cold water. Acetylsalicylic acid (0.51 grams) was also dissolved in 20 mL of 95%
ethanol. These were combined, stirred, and set aside for about 15 minutes until
copper (II) acetate crystals formed. Melting point analysis was performed next
on only the acetylsalicylic acid sample and some salicylic acid. Before
performing the experiment, the literature melting points were researched and
recorded for reference. A very small amount of aspirin and salicylic acid were
obtained and placed in two tiny rubber-stopped tubes. These were fastened to a
thermometer, and the complex was submerged in a test tube of oil. The test tube
was secured by a clamp and metal block complex to a hot plate. As the
temperature rose approached the anticipated melting points of both substances,
the powders (solids) turned to clear liquid. The temperature ranges in which
this occurred were recorded and compared with literature values. Infrared
spectroscopy was then performed on aspirin and copper (II) aspirinate. First,
two NaCl plates were washed with acetone and set aside to dry. A tiny sample of
aspirin was placed in a mortar containing several drops of Nujol. The pestle was
used to grind and mix these completely, and part of the mixture was smeared onto
one of the NaCl plates. These were sandwiched together, and placed into a slot
in front of the IR beam. The absorbance was then measured and recorded. This
process was repeated for copper (II) aspirinate. Ultraviolet (UV) analysis was
also conducted on copper (II) aspirinate. Samples were placed in three 1-cm
cuvettes for experimentation. The two control samples were water and 0.01 grams
of copper (II) acetate dissolved in water. A small amount of copper (II)
aspirinate crystals were dissolved in a third cuvette containing equal amounts
of water and 95% ethanol. These were then placed in the UV wells and their
absorbances at 800 nm were graphed and recorded. After synthesis, a pH titration
was performed on a 0.25-gram sample of acetylsalicylic acid in a 400-mL beaker.
After the aspirin was dissolved in about 300 mL of water, a pH electrode was
calibrated and inserted into the beaker with a magnetic stirring rod. A 25-mL
buret was cleaned and filled with a standardized 0.1030 M solution of NaOH. This
was then added at 0.5-mL intervals until within 2-3 mL of the anticipated
endpoint. Then the intervals were decreased to 0.2 mL and 0.1 mL of base. After
the titration was completed, the data was entered into SigmaPlot, and the
equivalence point was determined. The final experiment performed was a
conductimetric titration of acetylsalicylic acid. A 0.2-gram sample was weighed
and dissolved in a 400-mL beaker, and titrated with the same standard 0.1030 M
NaOH at 0.5-mL intervals. A magnetic stirring rod and a conductance electrode
were placed into the acidic solution, and conductance readings were recorded
after every 0.5 mL of base were added. This data was also graphed, and the
equivalence point was used to calculate the Ka for acetylsalicylic acid. Results
Aspirin Synthesis Theoretical yield (g): 3.95 g Actual yield (g): 3.71 g Percent
Yield (%): 94% Catalyst Effects Catalyst Time for 4º C change (min:sec)
Anhydrous Sodium Acetate 3:40 Pyridine 0:43 H2SO4 0:04 Copper (II) Aspirinate
Synthesis Theoretical yield (g): 0.61 g Actual yield (g): 0.48 g Percent Yield
(%): 79% Melting Point Analysis Acetylsalicylic acid (literature)3: 136º C
Salicylic acid (literature)4: 159º C Acetylsalicylic acid (experimental):
130º-132º C Salicylic acid (experimental): 160º-162º C Conductimetric
Titration of Aspirin Weight of aspirin sample (g): 0.2003 g Molarity of NaOH
(n/L): 0.1030 M Equivalence point (mL of NaOH): 11.51 mL Moles of NaOH at
equivalence point (n): 0.001186 n Molecular weight of aspirin (g/n): 169.0 g/n
Expected molecular weight (g/n): 180.2 g/n Percent error (%): 6.215% pH
Titration of Acetylsalicylic Acid Volume of NaOH at endpoint (mL): 14.69 mL pH
at halfway point (pKa): 3.53 Ka for acetylsalicylic acid: 3.0 x 10-4 Literature
Ka for acetylsalicylic acid5: 3.0 x 10-4 Error (%): 0.00% Infrared (IR)
Spectroscopy Peaks Nujol Expected6 (cm-1) 2900 1460 1377 723 Nujol Actual (cm-1)
2900 1460 1376 723 Aspirin Expected5 (cm-1) 1752 1688 1187 917 Aspirin Actual
(cm-1) 1754 1688 1187 917 Cu-Asp Expected5 (cm-1) 1756 1725 1692 1187 917 Cu-Asp
Actual (cm-1) 1603 1460 1377 1155 967 (Also see attached: IR and uv-vis spectra,
pH titration curve, and conductimetric titration curve.) Discussion The
underlying purpose of this experiment was to confirm that aspirin and copper
(II) aspirinate were indeed synthesized. The series of tests performed on the
products of the syntheses (spectra, titration, melting point) used the various
physical properties of the two crystals to do exactly this. Some data, due to
random error, deviated slightly from the expected experiment results. In the
synthesis portion of the experiment, the percent yield for aspirin was over 90%,
more than enough to carry out the copper (II) aspirinate synthesis and other
procedures. The yield for copper (II) aspirinate synthesis was not as high, but
since only small amounts of the blue crystal were needed for subsequent
analyses, this percent yield was acceptable. An excess of copper was present in
the dissolved copper acetate, so the limiting reagent in this reaction was
aspirin. The observed melting point for salicylic acid was extremely accurate
according to its literature (expected) value. However, the melting point for
aspirin appeared slightly lower than normal by about 5º C. This small error
most likely occurred because of the incomplete evaporation of the aspirin
crystals used in the analysis. The remaining water molecules in the crystals
probably acted as heat conductors, and therefore the temperature recorded on the
thermometer may not have reflected the temperature inside the small tube. The
presence of excess H2O in the acetylsalicylic acid sample was also observed in
the IR spectrum in the form of a premature dip in the anticipated straight
horizontal line between wavenumbers 3700 and 3400 (approximately). The product
apparently had not been completely evaporated in the oven prior to examination.
However, expected peaks for aspirin matched almost exactly those observed. Yet
due to the wetness of the product, the copper (II) aspirinate peaks differed
slightly from those expected, but accuracy was approximately 90%. The UV-vis
spectrum of copper aspirinate was somewhat rough because of incomplete
dissolution of the crystals in the alcohol/water mixture. At a glance, the two
graphs (copper aspirinate and copper acetate) appeared similar, but further
analysis showed that they were very different indeed. From 400-800 nm, the
absorbance of copper aspirinate oscillated from 0.355 to 0.325 to 0.363 and then
traveled back down. Copper acetate's absorbance remained relatively constant at
approximately 0.02 from 400-530 nm, at which point it rose rapidly and steadily
to a 0.48 peak at 760 nm. It was therefore concluded that the two crystals were
completely different, although they appeared similar in color to the naked eye.
In the pH titration of acetylsalicylic acid, the endpoint was determined and
used to calculate Ka. The result (measured to only two significant figures)
exactly matched the expected Ka of 3.0 x 10-4. This confirmed further that the
compound synthesized was indeed aspirin. This procedure was also performed after
observing the errant data due to excess water molecules in the sample, so the
crystals were heated again in an oven to completely evaporate them for further
use. If the sample had not been evaporated, fewer moles of actual acid would
have been dissolved, and the endpoint would have been significantly lower. The
resultant Ka would therefore have been much higher than expected. Further
evidence of the complete evaporation of the aspirin product was observed in the
final test, conductimetric titration. The equivalence point was measured and the
molecular weight was calculated within 10% of the expected value. The bulk of
the error in this titration was attributed to the erratic movements of the
stirring rod, which could have unequally distributed the incoming NaOH, thus
slightly altering the conductance. This error, however, was random, and nearly
negligible. References "Laboratory Handbook for General Chemistry II";
Eckerd: St. Petersburg, 2003; p. 33. "pH Titration of Acetylsalicylic
Acid." Eckerd College Intranet. 29 Apr 2003.