Introduction: ommercially prepared aspirin tablets are not considered 100% pure acetylsalicylic acid. Most aspirin tablets contain a small amount of binder which helps prevent the tablets from crumbling. The binder is chemically inert and was intentionally added by the manufacturer, but its presence means that aspirin tablets do not have 100 percent purity. Moreover, moisture can hydrolyze acetylsalicylic acid. Thus, aspirin which is not kept dry can decompose. Acetic acid is the hydrolysis product formed by the reaction of water with acetylsalicylic acid. You may have noticed the smell of vinegar (acetic acid) when opening an old bottle of aspirin, or a bottle which has not been properly sealed. The hydrolysis product responsible for this odor is acetic acid and is formed in the following way: For this lab, you will determine the purity of the aspirin by titration and by spectrophotometric analysis. Titration: You will determine the percentage of acetylsalicylic acid in your aspirin sample by means of titration. In this procedure, you will incorporate a process known as a back titration. In a normal titration, an experimenter is able to determine the amount of analyte present in a solution by carefully adding incremental volumes of a standard solution until reaction between analyte and titrant is judged to be complete. ccasionally, it is convenient or necessary to add an excess of the titrant and then titrate the excess with another reagent. This process is called back titration. In this technique, a measured amount of the reagent, which would normally be the titrant, is added to the analyte sample so there is a slight excess of reagent present. After this reagent reacts completely with the analyte, the amount of excess (unreacted) reagent is determined by titration with another standard solution. The amount of reagent which reacted with the analyte can be found by determining the difference between the amount of reagent added and the amount in excess: moles reagent reacted (with analyte) = total moles added excess moles backtitrated nce this quantity has been determined, stoichiometric considerations will allow you to find the moles of analyte initially present. At room temperature, acetylsalicylic acid can be neutralized with base. Page 1
If acetylsalicylic acid were the only acid present in your sample, you could determine the purity of the aspirin by a simple titration of your sample with sodium hydroxide. However, if acid impurities are present, titration of the aspirin will neutralize not only the acetylsalicylic acid, but the acid impurities as well. Thus, from such a titration, one could calculate the total number of moles of acid present in the sample by measuring the volume of standardized NaH required to reach the end point. If the stoichiometric ratios between acid and base are 1:1 (as in this experiment), the total number of moles of acid may be calculated by: moles acid = moles base = (liters base) x (molarity base) The titration of an impure sample of aspirin will yield the conjugate bases acetate ion, salicylate ion, and acetylsalicylate ion. f these, only the acetylsalicylate ion is an ester. It will react with additional base reasonably rapidly at elevated temperatures. This reaction represents what is termed a basepromoted hydrolysis, or saponification, of esters. After you have neutralized all acidic material in the aspirin by titration with base, you will add a known excess amount of base to cause the saponification to occur. The excess base that is not consumed in the hydrolysis will be determined by a backtitration with standard Hl. From your data, you will be able to calculate the grams of acetylsalicylic acid in your aspirin sample. The following example may help you understand this calculation: Example: A 0.5130g sample of aspirin prepared by a student required 27.98 ml of 0.1000M NaH for neutralization. An additional 42.78 ml of 0.1000 M NaH was added, and the sample was heated to hydrolyze the acetylsalicylic acid. After the reaction mixture cooled, the excess base was backtitrated with 14.29 ml of 0.1056 M Hl. How many grams of acetylsalicylic acid are in the sample? What is the percentage of acetylsalicylic acid (or the purity)? Solution: First recognize that the 27.98 ml of base was used to neutralize all acidic material present in the sample. Since we are only interested in the quantity of acetylsalicylic acid, we must determine the quantity of base required for hydrolysis of the ester. The total number of moles of base added for the hydrolysis reaction is moles NaH = 0.04278 L x 0.1000 M = 4.278 x 10 3 mol The number of moles of Hl used in the titration corresponds to the excess NaH, or the number of moles not consumed in the hydrolysis reaction: mol Hl = mol excess NaH = 0.1056 M x 0.01429 L = 1.509 x 10 3 mol The difference between the number of moles of base added for the hydrolyses and those which were not consumed equals the number of moles of base that brought about hydrolysis. This is Page 2
exactly equal to the number of moles of acetylsalicylate ion which is equal to the number of moles of acetylsalicylic acid: 4.278 x 10 3 1.509 x 10 3 = 2.769 x 10 3 mol The number of grams acetylsalicylic acid is found using the molecular weight of acetylsalicylic acid: grams = 2.769 x 103 mol x 180.2 g/mol = 0.4989 g acetylsalicylic acid Thus, % purity = 0.4989 g/0.5130 g x 100 = 97.25 % Spectrophotometric Analysis of Aspirin: A colored complex is formed between aspirin and the iron (III) ion. The intensity of the color is directly related to the concentration of aspirin present; therefore, spectrophotometric analysis can be used. A series of solutions with different aspirin concentrations will be prepared and complexed. The absorbance of each solution will be measured and a calibration curve will be constructed. Using the standard curve, the amount of aspirin in a commercial aspirin product can be determined. The complex is formed by reacting the aspirin with sodium hydroxide to form the salicylate dianion. H H 3 (s) + 3H (aq) (aq) + H (aq) + 2H (l) 3 2 The addition of acidified iron (III) ion produces the violet tetraaquosalicylatroiron (III) complex. +3 + [Fe(H ) ] 2 6 Fe(H 2) 4 + + H + H 2 3 + Note: The ph of the solution in which the complexation reaction is run must be kept between 0.5 and 2.0. If the ph gets too high the iron will precipitate out as a hydroxide and if the ph is too low the salicylate dianions will be protonated rendering them unavailable to the complexation reaction. The ph is controlled with hydrochloric acid. Page 3
Materials Needed: Burets (2), 600 ml beaker, 250 ml Erlenmeyer flasks (3), hot plate, IE 125mL Erlenmeyer flasks, 10 ml graduated cylinder, 5mL pipet, cuvettes, 250 ml volumetric flask, 25 or 50 ml volumetric flask (3 or more) hemicals Needed: Aspirin (nonbuffered commercial tablets), phenolphthalein solution, 95% ethyl alcohol, 0.1M HL 0.1M NaH (standardized), 1M NaH (not standardized), 0.025 M Fel 3 buffer solution (6.8 g Fel 3 6H 2 to 100 ml deionized water in a 250 ml beaker. Add 3.0 ml of concentrated Hl and 12.0 g of Kl. Dissolve and dilute to 1.0 L with deionized water, check ph) Page 4
Procedure Volumetric Analysis of Aspirin Prepare an ice bath. btain 75 ml ethyl alcohol (EtH) in a clean and dry 125mL Erlenmeyer flask. ool the EtH in the ice bath. btain two 25mL burets. Fill one buret with acid (Hl) and one with base (NaH). nce the burets have been filled, be sure to record the concentrations of the acid and base in your laboratory notebook. Remember, the buret reading can be estimated to ±0.01mL Weigh to the nearest milligram about 0.5g of aspirin tablets into a clean dry 250 Erlenmeyer flask. Add about 25mL of 95% ethyl alcohol that has been cooled to the flask and swirl the flask to dissolve the aspirin. Swirl to dissolve. Add 23 drops phenolphthalein and rapidly titrate with standard (~0.1 M) NaH, being sure to record starting and finishing volumes. The first 10 ml can be added quickly, followed by slow additions of NaH until a faint pink color persists. (This indicates the phenolphthalein end point.) The volume of NaH used for this titration corresponds to that which is required to neutralize all acids present in your sample, that is, impurities as well as the acetylsalicylic acid. Saponification of Aspirin: To saponify (or hydrolyze) the aspirin, you will add additional NaH from the buret, keeping careful record of the quantity used. The quantity of base to be used is about 15mL more than the volume of base used in the previous titration. It will probably be necessary to add more NaH to your buret. Record the initial volume, add the base to the Erlenmeyer flask, being careful to stop before reaching the lowest graduation mark (25.00mL) on the buret. Record the final volume. The volume of base added is the difference between these two readings. If the pink color should disappear, add 2 more drops of indicator. Backtitration of excess base: Record the initial volume of Hl in the second buret. Backtitrate the excess NaH in the Erlenmeyer flask with the standard (0.1 M) Hl solution until the pink color disappears. Record the final volume of Hl used. Repeat this procedure with your other two samples. alculations alculate the grams of acetylsalicylic acid in each of your aspirin samples and the percentage purity of the aspirin samples (see example). alculate the mean percentage purity and the standard deviation. Procedure for Spectrophotometric Analysis of Aspirin Preparing the Standards and Samples 1. Weigh 400 mg of acetylsalicylic acid in a 125 ml Erlenmeyer flask. Add 10 ml of a 1 M NaH solution to the flask and heat to boiling. Boil for five minutes and avoid splattering, washing down the sides of the flask with deionized water as needed. ool and quantitatively transfer this solution to a 250mL volumetric flask, and dilute to the mark with DI water. = Stock Solution 2. Pipet a 5.0 ml sample of this aspirin standard solution to a 50 volumetric flask (if using a 25mL flask pipet 2.5mL). Dilute to the mark with the 0.02 M iron (III) buffer solution. Label this solution "A". Page 5
3. Prepare similar solutions with 4.0, 3.0, and 2.0mL portions of the aspirin standard. Label these "B,, and D." Note: if using the 25 ml flask, use 2.0, 1.5, and 1.0 ml portions. 4. Using a commercial aspirin product, Place one aspirin tablet in a 125 ml Erlenmeyer flask. Add 10 ml of a 1 M NaH solution to the flask, and heat until the contents begin to boil. Quantitatively transfer the solution to a 250 ml volumetric flask, and dilute with distilled water to the mark. Pipet a 2.5 ml sample of this aspirin tablet solution to a 50 ml volumetric flask. Dilute to the mark with a 0.02 M iron (III) solution. Label this solution " unknown solution," and place it in a 125 ml Erlenmeyer flask. (Note: use 1.25mL when using the 25mL volumetric flask). Analysis 5. Turn on the spectrophotometer. Press the A/T/ button on the Spec 20 Genesys to select absorbance. Set the spectrophotometer to 530 nm. 6. Fill the one of the cuvettes about 3/4 full with 0.02 M iron (III) buffer. This is the "blank cuvette". 7. Fill one of the other cuvettes about 3/4 full with solution A. 8. Place the blank cuvette into the sample compartment of the spectrophotometer. Note: Before inserting a cuvette into the spectrophotometer, wipe it clean and dry with a kimwipe, and make sure that the solution is free of bubbles. Do not touch the clear sides of the cuvette. 9. Press 0 ABS 100%T to set the blank to 0 absorbance.. 10. Remove the blank cuvette from the instrument. 11. Place the cuvette containing solution A into the spectrophotometer. 12. Record the absorbance of solution A in your laboratory notebook. 13. Remove the cuvette containing solution A from the instrument. 14. Repeat steps 11 13 for solutions B,, and D, and the unknown solution. 15. Plot the graph of absorbance (yaxis) of the standards versus the concentration (xaxis) 16. Determine the concentration of the unknown from the graph using the absorbance measured. References John H. Nelson and Kenneth. Kemp, hemistry: The entral Science, Laboratory Experiments, 7/e, PrenticeHall, 1997. Daniel. Harris, "Quantitative hemical Analysis," 2 nd ed., W.H. Freeman, New York, 1987. Page 6