In the Laboratory edited by
Second-Year and AP Chemistry
John Fischer Ashwaubenon High School Green Bay, WI 54303-5093
Experiments with Aspirin
W
Londa L. Borer* Department of Chemistry, California State University, Sacramento, Sacramento, CA 95819;
[email protected] Edward Barry El Camino High School, Sacramento, CA 95821
Aspirin is much more than just a pain remedy. Aspirin is a compound that can be used to demonstrate many chemical principles in the high school classroom. In this series of experiments, we demonstrate how aspirin can be synthesized and characterized by methods that are easily accessible to high school students, how the hydrolysis of aspirin can be used as an introduction to kinetics, and how coordination chemistry (chelation) can be introduced by preparing and characterizing the copper complexes of aspirin and salicylic acid. Experiments Synthesis, Purification, and Characterization of Acetylsalicylic Acid Aspirin can be prepared by combining acetic acid and salicylic acid (SA) (which reacts at the phenol functional group) to form the ester. However, a more effective synthesis involves combining acetic anhydride with salicylic acid in the presence of a catalyst (Scheme I) (1–3). O
O
OH C
O
OH
+
OH C
CH3
O O
O
C
CH3
H
O
+ CH3
Acetic anhydride
CH3
O
where k is the rate constant dependent upon T, P, concentrations, ionic strength, etc., and [A] is the concentration of the reactant. ASA hydrolyzes in water to SA and acetic acid as follows: OH
O
C O
C
OH C
O
+
H OH
H2O
CH3
Acetylsalicylic acid (aspirin)
Acetic acid (vinegar)
Scheme I
A comparison of the effectiveness of a catalyst is made by having the students prepare aspirin using catalytic amounts of either concentrated H2SO4 or H3PO4. They report percent yield before and after recrystallization from ethyl alcohol and purity of sample determined from melting point and thinlayer chromatography.
The Percentage Composition of an Aspirin Tablet In most medicines today, the active ingredient is only a part of the capsule or dosage. Pharmaceutical companies add fillers to counter the side effects of the drug. Acetylsalicylic acid (ASA) is a strong acid with a pH ≈ 3. A coating is added to protect the lining of the throat and stomach. In this experiment, the filler is removed and the amount of ASA is determined. Four procedures for analyzing an aspirin tablet can be performed and are presented in this experiment. They are (i) analysis of the theoretical percent ASA in an aspirin tablet, (ii) analysis of ASA by weight, (iii) analysis of ASA by titration, and (iv) analysis of the filler. The experiments allow application of mass relationships, molar calculations, etc., as well as different experimental techniques of analysis and recovery of product. Determination 354
Kinetics of Hydrolysis of ASA The hydrolysis of ASA to SA has been the subject of many investigations and has been used as a model for investigation of mechanisms. Many kinetic runs are required to establish the dependence of “k”, the rate constant, on pH, temperature, buffer concentration, and ionic strength (4–6 ). The objective of the following experiments is to observe the effects of some of the variables on the pseudo-first-order reaction of the hydrolysis of ASA. The general rate law for first order is ᎑d[A]/dt = k[A] (1)
O
C
O Salicylic acid
of the melting point of the white solids collected will convince students of where they have collected ASA and where they have collected the filler, starch.
Acetylsalicylic acid (aspirin)
O
+
C
CH3
O
Salicylic acid
(2)
Acetic acid (vinegar)
To study such a reaction, a method that itself does not alter the course of the reaction must be chosen. Most studies use UV spectral analysis and monitor the decrease in ASA or the increase in SA in the UV region of the spectrum (≅ 290 nm). However, this is difficult in a high school laboratory because most high school labs do not have access to a UV spectrometer. However, they may have access to a visible spectrometer such as a Spectronic 20. We have chosen to follow the reaction by the appearance of the complexation of SA with ferric chloride, FeCl3. Ferric chloride reacts with the phenolic hydroxyl group of the SA produced in the hydrolysis of ASA to produce a purple iron complex. A 1:1 ratio of Fe3+ to SA is assumed in the formation of the complex. The concentration is measured as a function of the light absorbance at a wavelength characteristic of the product (7–9). For this study, three variables were chosen: concentration of ASA, solution pH, and concentration of the ferric chloride. In order that temperature will not be a variable, a constanttemperature bath is used. The students will first determine the peak absorbance for the Fe(III) complex of SA. Then they will compare the effect of FeCl3 on ASA and SA. Once they determine that Fe(III) reacts only to form a SA complex, they will determine the effect of pH on the formation of this
Journal of Chemical Education • Vol. 77 No. 3 March 2000 • JChemEd.chem.wisc.edu
In the Laboratory
complex. To do kinetics, they will determine the effect of concentration of ASA and the amount of FeCl3 present on the formation of the iron–salicylic acid complex. They will determine the effect of when FeCl3 is added on the formation of the iron–salicylic acid complex. Once all data are collected, graphs are prepared to show results. From the graphs, students are asked to evaluate the rate expression for the hydrolysis of aspirin in water from the rate expression Rate = k[ASA]m (3)
Synthesis of Copper(II) Aspirinate and Salicylate Aspirin is an NSAID, nonsteroidal antiinflammatory drug. There are many other NSAIDs, some of which are also used as pain relievers, but aspirin is the best known. Though scientists do not know for sure, they believe that all NSAIDs work in a similar way. Aspirin is thought to work by blocking the formation of certain prostaglandins (10). Prostaglandins play a role in inflammation, blood clotting, and regulation of immune response and neurotransmission. When the body suffers a trauma, normal tissue is disrupted and apparently prostaglandins are released as a cleanup crew. Inflammation follows, which may aggravate sensations of pain. Aspirin slows the entry of more prostaglandins into the area, thereby lessening the pain (11). There is no evidence that complexes of SA with Fe3+ are involved in blocking the release of prostaglandins. However, some scientists have proposed that complexation with Cu2+ is important in NSAID activity. It is known that aspirin complexes to metals, yet its effect on trace metals of the body are poorly characterized and understood. The role of metal complexes in biological systems is relatively unexplored and much work is necessary for a comprehensive understanding (12–14). In this experiment, synthesis of copper(II) aspirinate, Cu(ASA)2 (15), and Cu(II) salicylate, Cu(SA)2, are carried out and a comparison made of the two complexes. The structure of Cu(SA)2 can be represented as follows: COO
OOC Cu
O
O
Copper (II) salicylate
The pKa value for the carboxylic hydrogen is 2.88, while that of the phenol is 13.6 (16 ). The protons are removed from the SA by a base such as sodium carbonate or the acetate ion of copper(II) acetate. The SA is bonded to the copper by coordinate covalent bonds through the oxygen atoms.
Reaction of Copper (II) Aspirinate in Aqueous Solution The object of this experiment is to qualitatively observe the change of copper(II) aspirinate in aqueous solution over time. The copper ion, like the Fe(III) ion, forms an acidic solution in water and will catalyze the hydrolysis. Note that a UV–vis spectrometer is necessary for this experiment. All changes will be observed in the ultraviolet region of the spectrum. Conclusions This series of laboratory activities is very effective as a unit of study for advanced or second-year AP high school
chemistry students. The value of this series of laboratory activities lies in the following features First, the laboratory activities require students to synthesize a product using a number of procedures and then analyze their results to decide which procedure produces the best yield and purest product. This is valuable in allowing the students to study the relationship between proposed experimental procedures and theoretical product quality and yield. Second, the laboratory activities are designed in such a way that the product of one activity is utilized in subsequent activities. This is useful in demonstrating the idea of continuity in scientific research. Third, the experiments using spectrometry give students access to typical instrumentation and allow them to observe the function of the instrument. A further benefit of these activities lies in requiring teams of individuals to complete different studies of variables. This allows the entire activity to be completed in a reasonable amount of time and promotes effective communication between the teams of students. This division of labor promotes the collaborative aspects of scientific research. A fourth benefit of this series of activities lies in requiring the students to prepare a comprehensive report of their research. In reflecting and completing a report of the entire research activity, students develop a deeper understanding of the requirements of dealing with a research problem and then reporting the results. WSupplemental
Material Supplemental material for this article is available in this issue of JCE Online. Literature Cited 1. Pavia, D. L.; Lampman, G. M.; Friz, G. S. Introduction to Organic Laboratory Techniques, 3rd ed.; Harcourt Brace: Orlando, FL, 1988. 2. Hassell, C. A.; Marshall, P.; Hill, J. W. Chemical Investigations for Changing Times, 7th ed.; Prentice Hall: Englewood Cliffs, NJ, 1995; pp 193–198. 3. Brown, D. A.; Friedman, L. B. J. Chem. Educ. 1973, 50, 214. 4. Edwards, L. J. Trans. Faraday Soc. 1950, 46, 723–735; 1952, 48, 696–699. Garrett, E. R. J. Am. Chem. Soc. 1957, 79, 3401. Fersht, A. R.; Kirby, A. J. J. Am. Chem. Soc. 1967, 89, 4857. Spancake, C. W.; Mitra, A. K.; Kielsig, D. O. Int. J. Pharmacol. 1991, 75, 231– 239. Solomon, D.; Hur, C.; Lee, A.; Smith, K. J. Chem. Educ. 1996, 73, 173. 5. Alibrandi, G.: Micali, N.; Trussi, S.; Villari, A. J. Pharm. Sci. 1996, 85, 1105. 6. Loudon, G. M. J. Chem. Educ. 1991, 68, 973. 7. Sawyer, D. A.; Heineman, W. R.; Beebe, J. M. Chemistry Experiments for Instrumental Methods, Wiley: New York, 1984, 205. 8. Street, K. W. J. Chem. Educ. 1988, 65, 914. 9. Soloway, S.; Wilen, S. H. Anal. Chem. 1952, 24, 979. 10. Wesp, E. F.; Brode, W. R. J. Am. Chem. Soc. 1934, 56, 1037. 11. Aspirin and Other Salicylates, 1st ed.; Vane, J. R.; Botting, R. M., Eds.; Chapman and Hall Medical: New York, 1992. 12. Liska, K. Drugs and the Human Body, 3rd ed.; Macmillan: New York, 1990; p 332. 13. Arena, G.; Kovu, G.; Williams, D. R. J. Inorg. Nucl. Chem. 1978, 40, 1221. 14. VanRinsvelt, H. A.; Sater, R.; Hurd, R. W.; Andres, J. M. Nucl. Instrum. Methods Phys. Rev. B 1985, B10-B11(2), 660. 15. Dudek, E. J. Chem. Educ. 1977, 54, 329. 16. L’Heureux, G. A.; Martell, A. E. J. Inorg. Nucl. Chem. 1966, 28, 481.
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