An ASA tablet (theres an unknown amount of ASA but its nominal ASA content is 325 mg was dissolved in 1.0 mol/L NaOH solution (after heating), following the instructions of this experiment. The resulting mixture was transferred quantitatively into a 250 mL (0.2500 L) volumetric flask and diluted to volume with distilled water. The resulting solution is the stock solution of the tablet unknown.

A 2.00 mL sample of the stock solution of the unknown was pipetted into a 50.00 mL volumetric flask and diluted to volume with 0.02 mol/L iron(III) chloride solution dissolved in 0.1 mol/L HCl solution. This diluted unknown solution (purple) contains the iron-salicylate complex of the tablet. Look at the reactions in the lab to recognize that the ratio of the complex formed to ASA is 1:1

Use the data from Table 1 to determine the molar absorptivity, , and the concentration of the dilute unknown solution. The absorbance of this unknown diluted solution was measured to be 0.311.

Calculate the concentration (mol/L) of the original (stock) concentrated solution.

Knowing that the molar mass of ASA is 180.157 g/mol and that the unknown tablet was dissolved in 250 mL originally, calculate the mg of aspirin present in the stock solution.

Knowing that you should have 325 mg of ASA present (theoretical amount), calculate percent error. Show all calculations and your final answer must have the correct # of significant digits.

Table 1: Absorbance determined as a function of concentration of the salicylate-Fe complex iron-salicylate complex (moles/L) 8.82x10-5 1.77x10-4 2.69x10-4 3.64x10-4 4.66x10-4 Absorbance 0.097 0.192 0.292 0.396 0.508