The Determination of the Percent of Salicylic Acid in Aspirin

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1 Salicylic acid The Determination of the Percent of Salicylic Acid in Aspirin Aspirin or acetylsalicylic acid, may be prepared by the reaction of salicylic acid and either acetic acid or acetic anhydride in the present of a catalyst according to the following reaction: Salicylic acid Acetic acid KJ/mol 2 Fe Fe( 2 ) 3 6 or Fe 3 (aq) 3 1 atalyst 2 2 Fe 2 2 Violet Fe 3 omplex Aspirin owever, few organic reactions go to completion or give quantitative yields of the intended product. The presence of either salicylic acid or acetic acid in your aspirin indicates an impure product caused by either an incomplete reaction or the presence of moisture, reversing the reaction. In this experiment, the amount of salicylic acid in an aspirin sample, if any, will be determined spectrophotometrically using iron (III) ions. Iron (III) ions react quantitatively (K f = 3 x ) with the phenol (-) group on salicylic acid s benzene ring forming a highly colored violet complex according to the reaction: Absorbance = molar absorptivity, ε x cell width, l x concentration Absorbance = a constant ( ε l) x concentration or simply, A = abc, where a and b are constants To determine the purity of your sample, solutions with a known concentration of salicylic acid are prepared and the intensity of the light absorbed by these solutions measured using a spectrophotometer. The amount of light absorbed is related to the solution s color intensity which is directly related to its concentration according to Beer s Law. The algebraic form of Beer s law is 2 Water The molar absorptivity or molar extinction coefficient,,, is a number related to the nature of the absorbing species at the temperature and wavelength studied. The more intense the coloration of the absorbing species, the larger the value of the molar absorptivity. For example, at 635 nm, an aqueous copper solution, u 2 (aq) has a molar absorptivity of 10 Lmol/cm. If the copper ion is complex with ammonia, the highly colored violet species, tetraaminecopper (II), u(n 3 ) 4 2, is formed. With this complex, the molar absorptivity increases to 30 Lmol/cm. The cell width is also important. The greater the light path, the more light that is absorbed by the solution. When

2 measuring the amount of light removed by a solution, the solution is placed in a perfectly square cell called a cuvette. They may be made from plastic, glass, and even quartz. Most have an inside dimension of exactly one centimeter. In a similar experiment to this experiment, a student dissolved 2.75 grams of nickel (II) sulfate hexahydrate, NiS 4 6 2, in water and diluted the resulting solution to exactly milliliters. The absorbance of the nickel (II) sulfate solution was compared to solutions with known concentration. The results are provided in the table below. Given this information, both the concentration of the nickel sulfate solution and the purity of the original compound can be determined. A graph of the absorbance of the known solutions vs their concentration is represented in the following table: oncentration Absorbance M M M M M unknown The slope of the Beer s law plot above was determined to be with an intercept of giving the equation for the line to be: absorbance = concentration substituting the absorbance of the unknown solution of 0.225, the concentration of the unknown would become M. Note that the form of the equation for the absorbance graph is y = mx b, the same as Beer s Law, where the absorbance, A =,R. Recall that, is the molar absorbance, R the cell width of 1 cm, and the molar concentration. As a result, the molar absorptivity or molar extinction coefficient,,, can be determined from the data obtained. That is, the slope (,R) is numerically equal to the molar absorptivity in Lmol/cm since R, the cell width, is one centimeter. According to the experiment above, the molar absorptivity for nickel (II) sulfate solutions is 1.19 Lmol/cm at the room temperature and wavelength of the experiment. Before you determine the amount of salicylic acid in your aspirin sample, you must determine the optimum wavelength of light for your analysis. The optimum wavelength is the wavelength of the light were there are few competing species and a maximum absorbance, know as lambda max, 8 max. The wavelength at the maximum absorbance gives the spectrophotometer the greatest sensitivity to the substance under study, and hence the least error. Moreover, at the maximum absorbance, the curve is almost flat minimizing dilution errors. Let s study the absorption 2

3 spectrum of an aqueous nickel solution below. Note that there are three useful wavelengths in the adjacent nickel spectrum; 392 nm, 672 nm, and 720 nm. Which one should you choose? While the 392 nm peak will certainly give the maximum absorbance, it occurs at the limit of the detector sensitivity and may not be reliable. That leaves 672 nm and 720 nm. Both are relatively flat maximums but the 720 nm peak gives a slightly higher absorbance and therefore, a greater sensitivity. Fortunately, your iron-salicylic acid complex study will not be as complex as the nickel spectrum. Materials: iron (III)/l/Kl color reagent solution standard salicylic acid solution wash bottle w/distilled water 50 ml volumetric flask 100 ml beaker funnel 3-1 ml x 0.01 ml pipets 1-2 ml pipet pipet pumps for your pipets 8 - small beakers or plastic cups omputer w/vernier s Logger Pro software USB cable Spectrophotometer 9 - polystyrene cuvettes cuvette holder ethanol 1 ml calibrated polyethylene transfer pipet marking pen spatula Procedure: 1. Number seven small beakers or plastic cups with tape or marking pen. Prepare the six standard solutions according to the following table. That is, add 2.00 ml of the iron (III)/l/Kl color reagent solution to each of six small beakers or plastic cups. Also add 2.00 ml of the iron solution to a seventh beaker or cup which will be used for your unknown aspirin sample. Next add the distilled water in the amount prescribed in the table below followed by the standard salicylic acid solution using separate 1.00 ml measuring pipets. arefully mix the solutions in the cups by swirling and add them to cuvettes keeping the solutions in order in the cuvette holder. 3

4 Trial Volume of standard 0.00 ml 0.20 ml 0.40 ml 0.60 ml 0.80 ml 1.00 ml Volume of water 1.00 ml 0.80 ml 0.60 ml 0.40 ml 0.20 ml 0.00 ml Volume of color reagent 2.00 ml 2.00 ml 2.00 ml 2.00 ml 2.00 ml 2.00 ml Mass of salicylic acid in mixture 0.00 mg mg mg mg mg mg 2. The original standard solution of salicylic acid contained 500 mg of salicylic acid in 250 milliliter of solution. 25 ml of this solution was then diluted to 100 ml in a volumetric flask. Given this information, complete the table above by calculating the mass of salicylic acid in each of your solutions with a known concentration in milligrams/milliliter (mg/ml). 3. Tare a piece of weighing paper on a balance. Weigh 0.25 gram of your aspirin onto the weighing paper recording its mass to the precision of your balance. Break up any large pieces to facilitate the solution process. Place a funnel into a 50 ml volumetric flask and quantitatively transfer the aspirin sample into the funnel. Rinse the solid from the funnel into the volumetric flask with a 1-ml polyethylene transfer pipet in one milliliter portions until a total of 10 ml of ethanol has been added. Swirl the mixture until the aspirin dissolves completely in the alcohol. The solution process can be accelerated by heating the flask by swirling the it in hot tap water. When all the aspirin has dissolved, add distilled water to the flask until the meniscus of the solution level just rests on the calibration scribe on the flask. Stopper, and mix the solution by inverting the flask roughly twenty times. Pipet 1.00 ml of your aspirin solution into the seventh beaker or cup containing 2.00 ml of the iron color reagent and swirl gently to mix the solutions. No additional water will be necessary in this trial. 4. Turn on your computer and load the Logger Pro software by clicking on the Logger Pro icon. onnect the spectrophotometer to an available USB port on your computer with the proper cable. After a few seconds, your computer will automatically identify the spectrophotometer. 5. The next step is to calibrate the spectrophotometer. Fill an empty cuvette ¾ full with distilled water, called a blank, and place it in the cuvette compartment in the spectrophotometer. To correctly use a cuvette, remember: All cuvettes should be wiped clean and dry on the outside with a tissue. Always position the cuvette with the clear sides or reference point facing the white reference mark on the top of the spectrophotometer. andle cuvettes only by the top edge of the ribbed sides. All solutions should be free of bubbles. 6. lick on the Experiment drop down menu, scroll down to 4

5 calibrate, select spectrometer 1, and wait for the spectrophotometer lamp to warm properly. Place a cuvette containing distilled water into the cuvette holder, click Finish alibration, and then K. The spectrophotometer is now calibrated for every wavelength within its range. 7. To record the entire spectrum of the iron color reagent solution at once, place the cuvette containing solution #1 into the cuvette holder, and click on the green ollect icon or hit the space bar. When you are satisfied with the iron spectrum, click Stop or again hit the space bar. Remove the iron color reagent solution cuvette and replace it with solution #6 containing one milliliter of standard salicylic acid and click ollect or hit the space bar. Again, click Stop or hit the space bar to stop data collection. Leave the cuvette containing solution #6 in the cuvette holder. 8. nce an absorption spectrum is obtained for both the iron chloride and iron-salicylic acid complex, the spectrophotometer must be reconfigured to read the absorbance vs concentration at maximum absorbance. lick on the onfigure Spectrometer tool bar icon,. An absorption spectrum of the solution in the cuvette will appear with lambda max identified. To perform a Beer s Law study, check the Abs vs concentration option, then Done. You will now be asked if you want to save this trial. lick Yes and remove solution #6 from the cuvette holder. 9. Start the data collection by either hitting the spacebar or clicking on the green ollect icon. You are now ready to collect absorbance data for the six standard salicylic acid solutions. Starting with the cuvette containing solution #1, wipe the outside of the cuvette with a tissue and place it in the cuvette holder. Wait for the absorbance value displayed on the monitor to stabilize. lick on the Keep icon, (next to the green collect icon), enter the number of milligrams of salicylic acid in the mixture, 0" for this trial, and press the Enter key. The data pair you just collected should now be plotted on the graph. 9. Repeat the this procedure with the remaining five standard salicylic acid test solutions. When you have finished with your test solutions, click on the red Stop icon. 10. Examine the graph of absorbance vs. concentration. To test if the curve represents a direct relationship between these two variables, click the Linear Fit icon,, located on the Logger Pro toolbar. A best-fit linear regression line will be shown for your data points. This line should pass near or through the data points and the origin of the graph. Record both the slope, m, and y-intercept, b. 11. The next step is to determine the absorbance value for your aspirin extract solution. Add 5

6 the mixture in beaker or cup #7 to a clean, dry cuvette. As with your standard solutions, wipe the outside of the cuvette and be certain the clear sides of the cuvette are facing the detector. Monitor the absorbance value displayed in the absorbance meter on your Logger Pro screen. When this value has stabilized, record this absorbance to the nearest (Important: The reading in the meter is live, so it is not necessary to click ollect to read.) 12. Spectrophotometric methods of analysis are most accurate between about 20% and 70% transmittance (absorbance values between ~0.15 and ~0.70). If the absorbance of your aspirin solution is below 0.15, your results can be improved by employing a technique called the Single-Point Standard Addition method. The principle of this technique is to add a know amount of salicylic acid to your unknown sample, known as a spike, and measure the absorbance of the spiked solution, raising its absorbance into the range where the spectrophotometer is most accurate. To this end, add 0.1 ml of the standard salicylic acid solution used to prepare your Beer s Law plot above to the cup or beaker used for your unknown aspirin solution. Pour the solution back and forth between the cuvette and mixing cup several times to assure a good mix. Refill the same cuvette, wipe the outside of the cuvette, and measure the absorbance of this solution. Record this value. 12. Exit Logger pro, turn off your computer, and disconnect the data collection system. Throughly rinse all materials and return them to their designated areas. alculations: Q1. According to Beer s Law, the absorbance of a solution varies directly with the concentration of the solution. Using the Beer s Law relationship, calculate the mass of salicylic acid per ml in your aspirin sample. Beer's Law: Absorbance = (a constant) x concentration = abc = ab (a constant) x c Notice that Beer s Law has the form: y = mx b, where y is the absorbance, m, the slope, is a constant, x is the concentration, and b represents of y intercept, or a slight coloration in the blank. Q2. If the initial absorbance of your aspirin sample was too low for an accurate impurity determination and you spiked the sample with additional salicylic acid, calculate the concentration of your unknown aspirin solution using the relationship: where: x = concentration of your unknown aspirin solution A x = the original absorbance of your aspirin solution s = concentration of the standard salicylic acid solution x AV x s s = (A V ) (A V ) spiked total x x 6

7 V s = the additional amount of standard salicylic solution added (0.10 ml) V total = total volume of solution after being spiked (3.10 ml) A spiked = the absorbance the spiked aspirin solution V x = volume of your aspirin solution in the cuvette (3.00 ml) Q3. (a) The aspirin solution in the cuvette was actually diluted by a factor of three. That is, 1.00 ml of aspirin was added to the cuvette along with 2.00 ml of the iron (III) reagent. Therefore, the concentration calculated in question #2 must be multiplied by three to obtain the actual concentration of salicylic acid in your aspirin sample. (b) alculate the percent of salicylic acid impurity in your sample from the mass of the aspirin added to the volumetric flask, the volume of the volumetric flask (50-ml), and the mass of salicylic acid per milliliter as determined by the absorbance data in this experiment. DATA: Mass of aspirin: gram Trial blank #1 #2 #3 #4 #5 mg of salicylic acid in the sample 0.00 mg mg mg mg mg mg Absorbance Linear regression relationship for your standardization curve: m = b = riginal absorbance of your aspirin solution: Absorbance of the spiked aspirin solution: oncentration of salicylic acid in the solution: mg/ml % salicylic acid impurity in your aspirin sample: % 7

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