1 A COLORIMETRIC DETERMINATION OF ASPIRIN IN COMMERCIAL PREPARATIONS ADDITIONAL READING The concepts in this experiment are also discussed in sections 4.5 and 6.1, 6.2 of Chemistry and Chemical Reactivity by Kotz, Treichel, Townsend, and Treichel and in sections 4.5 and 6.3 of Mindtap – General Chemistry , by Vining, Young, Day, and Botch. Also, students are strongly encouraged to review the background information for the Beer’s Law experiment that was done in CHEM 1033 (Gen Chem I lab). This information is included at the end of this document. ABSTRACT This experiment is divided into two parts. In Part A, a known sample of acetylsalicylic acid, commonly known as aspirin, was treated with sodium hydroxide to form a known concentration of salicylate ion. You will react salicylate ion with the iron(III) cation to form a brightly colored solution (the color is due to the salicylato iron(III) complex ion that will be discussed in the pre-lab lecture). You will prepare various known concentrations of this colored solution and measure the absorbance of each one. Using Excel, you will make a Beer’s Law plot (also know as a calibration curve) of absorbance (y-axis) vs. concentration (x-axis), and use Excel to fit a straight line to the data and calculate the equation of the line. In Part B you will react an aspirin tablet with sodium hydroxide and then with iron(III), as in Part A, and measure the absorbance of the solution. Using your Beer’s Law plot from Part A, you will be able calculate the concentration of the salicylate ion and the mass of the acetylsalicylic acid in the aspirin tablet, and compare this mass will that printed on the bottle of aspirin tablets. BACKGROUND
2 Colorimetry, sometimes called visible absorption spectrometry, is a very important quantitative technique that is often used to determine concentration. If one looks at Beer’s Law, it may appear that you only need to measure the
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- Spring '10