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 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 4 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) 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 138 and 140 C. The melting point range for salicylic acid is 158-160 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+ 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 4 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 Bchner 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 4 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: 136 C Salicylic acid (literature)4: 159 C Acetylsalicylic acid (experimental): 130-132 C Salicylic acid (experimental): 160-162 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 5 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. "Acetylsalicylic Acid." ChemicalLand21.com. 30 Apr 2003. "Salicylic Acid." Wikipedia. 30 Apr 2003. McMurry, J.; Fay, R. C. "Chemistry."; Prentice-Hall: Upper Saddle River, 2001; 3rd ed. p. A-15. "Spectroscopy for the Aspirin Project." Eckerd College Intranet. 29 Apr 2003. Appendix A: Percent Yield Calculations (1.01 g sal / 138.1 g sal/n) x 180.2 g asp/n = 1.32 g asp (1.00 g sal / 138.1 g sal/n) x 180.2 g asp/n = 1.30 g asp (1.02 g sal / 138.1 g sal/n) x 180.2 g asp/n = 1.33 g asp 1.32 g + 1.30 g + 1.33 g = 3.95 g Asp (theoretical yield) Cu2Asp4 limiting reagent: (0.33 g Cu / 63.546 g/n) = 0.0052 n Cu2+ (0.51 g Asp / 169.0 g/n) = 0.0030 n Asp (0.0030 n Asp / 4n)(803.092 g/n) = 0.61 g Cu2Asp4. Appendix B: Conductimetric Titration Calculations (0.01151 L NaOH) x (0.1030 n/L) = 0.001186 n NaOH = 0.001186 n Asp. (0.2003 g Asp)/(0.001186 n Asp) = 169.0 g/n. (actual) (12.011 x 9) + (1.00794 x 8) + (15.9994 x 4) = 180.2 g/n. (theoretical) Appendix C: pH Titration of Aspirin (14.69 mL)/2 = 7.35 mL, which corresponds to a pH of approximately 3.53. 10-3.53 = 2.951 x 10-4 ~ 3.0 x 10-4. Appendix D: Spectroscopy Accuracy Calculations Copper (II) Aspirinate: (1756 + 1725 + 1692 + 1187 + 917) = 7727 (expected) (1603 + 1460 + 1377 + 1155 + 967) = 6562 (actual) 6562 / 7277 x 100% = 90.17% accuracy Aspirin: (1752 + 1688 + 1187 + 917) = 5544 (expected) (1754 + 1688 + 1187 + 917) = 5546 (actual) (5544 / 5546 x 100%) - 100% = 0.04% error