The purpose of this experiment is to evaluate the percent aspirin on a commercial aspirin tablet using an instrumental method, spectrophotometry. In a spectrophotometer, light from a strong lamp passes through a monochromator, which breaks the light into its component colors using a grating, then uses mirrors and slits to direct a light beam of the desired wavelength through the sample compartment, where you place a tube (cuvette) of solution. A detector on the other side of the sample compartment measures the amount of light that passes through the solution and compares that to the amount of light that passes through a control solution.
The result is displayed as either transmittance or absorbance. Transmittance is the ratio of light passing through the sample (I) to light passing through the control (I0). Absorbance is the negative logarithm of transmittance: T? I I 0 A?? LOGT?? LOG I I 0 The absorbance of a solution depends on three things: the nature of the solution, the distance the light travels through the solution (called the path length), and the concentration of the solution. This is summarized by the Beer-Lambert Law, A = abc A is the measured absorbance a is a constant (the absorptivity); it depends on ?and the nature of the solution b is the path length c is the solution concentration,
If the type of material and the cuvette diameter are always the same, ab is constant and the measured absorbance depends directly on the concentration of the solution. The absorbance is most sensitive to concentration if measured at a wavelength of high absorbance for that material. To use the spectrophotometer, turn it on and let it warm up for 5 minutes, so the lamp and detector outputs are stable. Select the desired wavelength, and use one knob (usually on the left) to set the detector to read A = ?or T = 0% when no light passes through the sample compartment (this is called the dark current).
Put a cuvette of control solution in the sample compartment, and adjust the other knob (usually on the right) to read A = 0 or T = 100%. The detector is now set for the range of light possible for your experiment. Put a cuvette of sample solution into the sample compartment and record the absorbance or transmittance of the sample. To prepare the aspirin, you will hydrolyze it with NaOH, then react it with acidic iron (III) ions. The Fe3+ ions and salicylate ions form a deep violet complex, which absorbs strongly at 530 nm.
Procedure Prepare iron (III)–salicylate complex solutions of known concentration (your standards) 1. Weigh about 0. 16 g pure acetylsalicylic acid into a clean 50 mL flask and record the exact mass used. Add 5 mL of 1 M NaOH solution and heat to boiling, being careful to avoid spattering. Allow the reaction mixture to cool, then transfer all of it into a 100 mL volumetric flask, using deionized water from your wash bottle to assure complete transfer. Add deionized water to bring the solution to exactly 100 mL. Mix thoroughly. This is your standard stock solution. 2. Label 5 small clean dry flasks A – E.
Use a 1–mL pipet to add 0. 50 mL of the standard stock solution into flask A, then use your 10–mL Mohr pipet to add 9. 50 mL acidified 0. 02 M FeCl3 solution, for a total of exactly 10. 00 mL. Mix thoroughly. Note the exact solution volumes used. 3. Prepare standards B – E as in step 2, using 0. 40 mL, 0. 30 mL, 0. 20 mL, and 0. 10 mL of standard stock solution and an appropriate amount of acidified 0. 02 M FeCl3 solution to give a total volume of 10. 00 mL. If you goof while pipetting and add a bit too much or too little, you don’t need to repeat that sample: just record the exact volumes used.
Prepare your aspirin sample 4. Perform the same hydrolysis as in step 1, using about 0. 16 g of ground aspirin tablet. Make one dilution of the sample stock solution, using 0. 30 mL of your sample stock solution + 9. 70 mL acidified 0. 02 M FeCl3 solution. Determine of absorbance of each solution 5. Set up the spectrophotometer at 530 nm, as described in the introduction. Use acidified 0. 02 M FeCl3 solution (the diluant) as the control solution. Read the absorbance of each standard A – E, and of the diluted aspirin sample (if your spectrophotometer has an analog readout, read %T and convert later).
You need not rinse the cuvette between samples (in fact, leftover water introduces more error than the same amount of leftover solution). Instead, use a disposable Beral pipet to withdraw as much solution as possible, then go on to the next solution. Analysis and Discussion Calculate the concentration of each standard, in grams acetylsalicylic acid per mL. Convert %T to absorbance, if necessary, then plot the absorbance of standards A – E as a function of concentration. Determine the slope and correlation coefficient of the graph. Mark the absorbance of the diluted sample on the standard curve and determine its concentration from the slope.
Calculate the concentration of aspirin in the sample stock solution, and compare that to the mass used. Calculate the % by mass of acetylsalicylic acid in your sample. As always, be attentive to significant figures. If you used a commercial aspirin tablet, calculate the % acetylsalicylic acid in an aspirin tablet from information given on the aspirin bottle (you may need to weigh a few aspirin tablets), and compare this calculated value to your results. Discuss the precision of your standard curve, including both the agreement between the points and the line and the value of the y intercept. Discuss the accuracy of your result.