Expt 1 - A Colorimetric Determination of Aspirin - SP 2015 - A COLORIMETRIC DETERMINATION OF ASPIRIN IN COMMERCIAL PREPARATIONS ADDITIONAL READING The

Expt 1 - A Colorimetric Determination of Aspirin - SP 2015...

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1 A COLORIMETRIC DETERMINATION OF ASPIRIN IN COMMERCIAL PREPARATIONS ADDITIONAL READING The concepts in this experiment are also discussed in sections 4.4 and 7.2 of Principles of Chemistry A Molecular Approach , by Tro. 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 available on Blackboard. ABSTRACT This experiment is divided into two parts. In Part A, a known sample of acetylsalicylic acid, commonly known as aspirin, is treated with sodium hydroxide to form the salicylate ion. This ion then reacts 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. A Beer’s Law plot (also know as a calibration curve) of absorbance (y-axis) vs. concentration (x-axis) will yield a straight line from which you will determine the slope 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 When a beam of light passes through a transparent medium (in this experiment a solution), its energy or intensity is reduced by several factors. Among these are 3 that are characteristic of the medium (solution). They are summarized in an expression known as Beer’s Law: bc I I H ¸ ¹ · ¨ © § 0 log where I 0 is the initial intensity of the light, I is the intensity after passing through the medium, ε is an absorption coefficient, commonly called the molar absorptivity (and also the extinction coefficient), that is characteristic of the absorbing species in the medium, b is the pathlength of the light beam in the medium in centimeters (and equal to the diameter of the cuvette used in this experiment), and c is the concentration of the absorbing species in mol/L. The ratio I / I 0 is also equal to the percentage transmittance, % T , divided by 100, i.e., % T /100, while the log of 100/% T is equal to the absorbance, A . Thus: ¸ ¹ · ¨ © § ¸ ¸ ¹ · ¨ ¨ © § 100 %T log I I log 0 A T ¸ ¹ · ¨ © § % 100 log ± ² T A % log 2 ³ In terms of absorbance, Beer’s Law is simply: A = ε bc . Both the absorbance and the percentage transmittance are dimensionless (unitless), hence ε has units of L/(mol.cm). 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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2 absorbance of a solution of unknown concentration, and then using Beer’s Law, you could use this measurement to determine the concentration (direct method).
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  • Spring '10
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  • Chemistry, Mole, Laboratory glassware, Salicylic acid, Acetylsalicylic acid

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