Spectrophotometry - Determining the Purity of Aspirin

Introduction:

In this experiment you will use a procedure based on the tendency of a complex derived from aspirin to absorb light. This method is called spectrophotometric analysis and the instrument used is a Spectrophotometer. In solution, aspirin itself does not absorb light in the visible range. However, when it is converted to an iron (III)-salicylate complex, it does absorb in the visible range. The absorption characteristics of this colored solution can be used to determine quantitatively how much aspirin is in the solution.  While this method of spectroscopic analysis has many important applications in both biology and chemistry, the basic principles of measuring a given property versus concentration are the basis of all analytical techniques.

Theory:

Under basic conditions, aspirin reacts with water (hydrolyzes) to form the salicylate dianion (has two separate negative charges) according the equation below:

Another product of this reaction is acetic acid (CH3CO2H) which is the active ingredient of vinegar.  So, you should detect a distinct odor of vinegar as the reaction proceeds.  You may also note this same odor when you open an old bottle of aspirin.  This is because this reaction occurs very slowly even at room temperature (particularly with lower quality generic brands).

When this salicylate dianion is mixed with an acidic solution of FeCl3(aq), iron(III) chloride, a highly colored (violet) iron (III)-salicylate complex is produced:

The violet color of the complex results from the fact that the complex strongly absorbs green light.  When this green is removed from normal white light, we observe violet (therefore, green is the compliment of violet).  This absorption of green light can be used to quantitatively determine the amount of aspirin present in the solution.  The more absorbance of green light, the more violet the solution, and hence, the more aspirin is present.  

If green-yellow light with a wavelength of 530 nanometer is directed into a solution that contains this aspirin complex, some of the green light will be absorbed:

As you can see, the intensity of the green light leaving the sample, I, is less than the original intensity of the green light, I0.  There are two ways of expressing this difference.  We can talk about the fraction of light that was transmitted through the sample,  transmittance (T); or we can talk about the amount of light that was absorbed by the sample, absorbance (A).  As you can see, one is opposite of the other:

transmittance (T)
absorbance (A)
T  =  I / I0
A  =  log (I0 / I)  =  log (1 / T)

The inverse relationship between transmittance and absorbance can best be seen in the following figure:


Notice that the %T can vary from 0 to 100% whereas the absorbance varies from 2.00 to 0.00 absorbance units.  The more light that passes through the sample,  the higher the transmittance and the lower the absorbance.  Conversely, the less light that passes through the sample, the lower the transmittance and the higher the  absorbance.

Unfortunately, a plot of transmittance versus concentration does not result in a straight line.  However, a plot of absorbance, versus concentration does provide a straight line:


In a typical experiment, several solutions of known concentration of the salicylate complex are prepared.  Since the concentration of these solutions is known, they are called standard solutions.  The absorbance of each standard solution is measured at the wavelength of maximum absorption (530 nanometer from the spectrum above) using a spectrophotometer. A graph of these absorbance values versus the concentration of each of the standards should yield a straight line. This relationship is known as Beers' Law::

A = a b c

In this equation, A is the absorbance of the solution, a is the molar absorptivity (a constant for this complex), b is the path length of cuvette (in cm), and c is the molar concentration of the solution being measured.  If the same cuvette is used to measure all of the solutions, then a and b are constant.  This means that the absorbance of a solution is directly proportional to the concentration of that solution.  Therefore, the molar concentration, c, of a solution can be determined by simply measuring the absorbance, A, of that solution.  Although we are actually measuring the absorbance of the complex, the stoichiometry of the reaction producing the complex is 1:1. So, if we know the concentration of the complex, we know the concentration of the aspirin is the same.

O.K., lets work through an example to see how all of this theory works.  Lets assume that 0.400 g of pure acetylsalicylic acid (aspirin) is treated as outlined in the experimental procedure below.  The concentration of the complex in the "Stock Solution" can be calculated as follows (remember, the molar mass of acetylsalicylic acid C9H8O4 is 180.2 g/mol):

0.400 g aspirin x (1 mol aspirin / 180.2 g)  =  2.22 x 10-3 mol of aspirin in the Erlenmeyer flask

Upon hydrolysis, and dilution to 500 mL (0.5 L), the molarity of the solution is:

M  =  2.22 x 10-3 mol  / 0.5 L  =  4.44 x 10-3 M

The "Stock Solution" is then diluted in varying proportions (aliquots) to yield the standard solutions "A", "B", "C", "D", and "E".  Solution "A" is produced by diluting 5 mL of the "Stock Solution" with 50 mL of FeCl3.  The concentration of aspirin in solution "A" can be found using the relationship:

M1V1 =  M2V2

where M1 is the molarity of the "Stock Solution", M2 is the molarity of the solution "A", V1 is the volume of the "Stock Solution", and V2 is the volume of the solution "A":

(5.0 mL) (4.44 x 10-3 M)  =  (50 mL) (M2)

Therefore, the concentration of standard "A" is 4.44 x 10-4 M.  Now that you know the concentration of standard "A", you can use the spectrophotometer to measure it's absorbance.  In this example, it had an absorbance of 0.50.  Likewise, you can determine the concentration and absorbance for each of the other standard solutions:

Solution
mL of Stock
Concentration
 Absorbance
"A"
5.0
4.44 x 10-4 M
 0.50
"B"
 4.0
 3.55 x 10-4 M 0.42
"C"
 3.0
 2.66 x 10-4 M 0.29
"D"
 2.0
 1.78 x 10-4 M 0.18
"E"
 1.0
 8.88 x 10-5 M 0.10

Now you have the data you need to create your Beers' Law plot.  However, it would be a good idea to check your data to make sure it is consistent before you throw away your "Stock Solution".  Remember, the whole idea behind this experiment is that the absorbance of a given solution will be directly proportional to the concentration of the aspirin in that solution.  If that is the case, then the Absorbance of a solution divided by the mL of Stock used to create it should be very nearly constant.  For example, if I divide the measured Absorbance of Solution "A" (0.50) by the mLs of Stock solution (5.0 mL), I obtain a value of 0.100 Absorbance/mL.  Likewise, I obtain values of 0.105, 0.097, 0.090, and 0.100 for solutions "B", "C", "D", and "E" respectively.  Since values are all within about 10% of each other, I am confident in the data I have collected and am ready to create my Beers' Law plot.  Remember that this is sample data that I have create to make the Beers' Law plot look good.  You may notice that the higher concentration solutions don't show as much Absorbance/mL as the lower concentration solutions.  This can happen if you use a large sample of aspirin.  If this happen, you will have to throw out the higher concentration result and only used the lower concentration results.

Once you have determined the concentration and absorbance for all five standards, you will plot these points using an 'X-Y Scatter' plot (Excel or Mr. Plot).  Your  Beers' Law plot should look like the one below:


Note that most of the points do not fall directly on the line.  So, we have asked the software to draw the 'best' straight line through the data.  This is the 'Least Squares Fit' or 'Trendline'.  The plot is fairly straight and has a 'goodness' of fit (R2) of 0.9934, where 1.000 is a perfect fit.  It also gives us an equation for the line which we will use to calculate the concentration of your aspirin.

Next you will need to process a sample of the aspirin you synthesized last week.  Lets assume that you used 0.400 g of your aspirin and processed it in exactly the same manner as you did the pure aspirin.  You will end up with 500.00 mL of a 4.44 x 10-3 M "My Aspirin" solution (assuming it is pure).  You then take 3.0 mL of this "My Aspirin" solution and diluted it to 50.0 mL with FeCl3.  The resulting solution has a concentration of 2.66 x 10-4 M (again, assuming it is pure).  You then measure its absorbance and obtain a value of 0.28. 

When you plotted your five standards, you obtained an equation for the linear regression equation.  In our example, that equation was:

Y = 1181·X - 0.014

In this equation, 'Y' is the absorbance, 'X' is the concentration of the solution, '1181.0' is the slope of the line, and '0.014' is the y-intercept.  Since we know the absorbance ('Y'), we can solve for the concentration ('X'):

X = (Y + 0.014) /  1181
X = (0.28 + 0.014) / 1181
X = 2.49 x 10-4 M

This is the actual concentration of pure aspirin in your sample.  However, we calculated that the sample should have a concentration of 2.66 x 10-4 M if it was truly pure.  Therefore, the purity of your aspirin is:

(2.49 x 10-4 M / 2.66 x10-4 M) x 100  =  93.6%

Not bad...but don't expect a call from Bayer anytime soon!

Procedure:

Preparation of Standards for the Beers' Law Plot:

In this section you will produce five acetylsalicylic acid standards of known concentrations.  Spectrophotometric determination of each standard's absorbance will be recorded and this data will be graphically plotted against concentrations to give a standard curve (Beers' Law Plot).

  1. Obtain a 125 mL Erlenmeyer flask and clean it thoroughly with soap and water.  Rinse it with dionized water and use a paper towel to remove as much of the excess water as possible.  Use your Bunsen burner to flame dry the Erlenmeyer flask.  BE CAREFUL NOT TO BURN YOURSELF.   Be sure to let the flask cool to room temperature and record its mass to the nearest 0.001 g.
  2. Add approximately 0.2 g of pure acetylsalicylic acid to the flask and record the mass to the nearest 0.001 g. By difference you can determine the mass of the pure acetylsalicylic acid that has been added.
  3. Add 5 mL of 1 M sodium hydroxide (NaOH) and heat to boiling (10-15 minutes).  Care should be exercised to avoid splattering and loss of contents, DO NOT allow all of the water to boil off.  CAUTION! NaOH is harmful to the skin and eyes. Rinse the inside walls of the flask with small portions (3-5 mL) of distilled water several times to ensure quantitative hydrolysis of the acetylsalicylic acid.
  4. Allow the solution to cool to room temperature.
  5. Quantitatively transfer the resulting solution of sodium salicylate to a clean 500 mL volumetric flask and then dilute with distilled water to the 500 mL mark.  Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label the flask as "STOCK SOLUTION."
  6. Using a 10-mL graduated pipette, transfer a 10.0 mL aliquot into a 50 mL volumetric flask and dilute to the 50 mL mark with 0.02 M FeCl3 solution.   Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label the flask as "Solution A".
  7. Rinse your cuvette with "Solution A" and then discard.  Refill the cuvette with "Solution A" and measure its absorbance.
  8. Using a 10-mL graduated pipette, transfer a 8.0 mL aliquot into a 50 mL volumetric flask and dilute to the 50 mL mark with 0.02 M FeCl3 solution.   Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label the flask as "Solution B".
  9. Rinse your cuvette with "Solution B" and then discard.  Refill the cuvette with "Solution B" and measure its absorbance.
  10. Using a 10-mL graduated pipette, transfer a 6.0 mL aliquot into a 50 mL volumetric flask and dilute to the 50 mL mark with 0.02 M FeCl3 solution.   Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label the flask as "Solution C".
  11. Rinse your cuvette with "Solution C" and then discard.  Refill the cuvette with "Solution C" and measure its absorbance.
  12. Using a 10-mL graduated pipette, transfer a 4.0 mL aliquot into a 50 mL volumetric flask and dilute to the 50 mL mark with 0.02 M FeCl3 solution.   Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label the flask as "Solution D".
  13. Rinse your cuvette with "Solution D" and then discard.  Refill the cuvette with "Solution D" and measure its absorbance.
  14. Using a 10-mL graduated pipette, transfer a 2.0 mL aliquot into a 50 mL volumetric flask and dilute to the 50 mL mark with 0.02 M FeCl3 solution.   Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label the flask as "Solution E".
  15. Rinse your cuvette with "Solution E" and then discard.  Refill the cuvette with "Solution E" and measure its absorbance.
  16. Check your data to make sure your absorbance data is decreasing relative to the decreasing concentration of each solution.  For example, the absorbance for the 4 mL solution should be half of that for the 8 mL solution and the absorbance for the 2 mL solution should be half of that for the 4 mL solution, etc.  If you find that the 10 mL solution shows significantly less absorbance that it should, it is possible that it is too concentrated and has fallen off the linear portion of the Beer's Law plot.  To correct this, you will have to run a sixth standard, "Solution F", using a 1.0 mL aliquot of the stock solution.  This will still give you 5 solutions to plot (Solutions B, C, D, E, and F) while allowing you to throw out Solution A (10 mL).
Determining the Purity of Your Aspirin:
  1. Obtain a 125 mL Erlenmeyer flask and clean it thoroughly with soap and water.  Rinse it with dionized water and use a paper towel to remove as much of the excess water as possible.  Use your Bunsen burner to flame dry the Erlenmeyer flask.  BE CAREFUL NOT TO BURN YOURSELF.   Be sure to let the flask cool to room temperature and record its mass to the nearest 0.001 g.
  2. Add approximately 0.2 g of the aspirin you synthesized in last week's lab to the flask and record the mass to the nearest 0.001 g.  By difference you can determine the amount of the aspirin that has been added.
  3. Add 5 mL of 1 M sodium hydroxide (NaOH) and heat to boiling (10-15 minutes).  Care should be exercised to avoid splattering and loss of contents, DO NOT allow all of the water to boil off.  CAUTION! NaOH is harmful to the skin and eyes. Rinse the inside walls of the flask with small portions (3-5 mL) of distilled water several times to ensure quantitative hydrolysis of the acetylsalicylic acid.
  4. Allow the solution to cool to room temperature.
  5. Quantitatively transfer the resulting solution of sodium salicylate to a clean 500 mL volumetric flask and then dilute with distilled water to the 500 mL mark. Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label this flask as "STOCK SOLUTION".
  6. Using a 10-mL graduated pipette, transfer a 5.0 mL aliquot into a 50 mL volumetric flask and dilute to the 50 mL mark with 0.02 M FeCl3 solution.   Be sure to thoroughly mix this solution by inverting the volumetric flask at least ten times.  Label this flask as "My Aspirin".
  7. Rinse your cuvette with "My Aspirin" and then discard.  Refill the cuvette with "My Aspirin" and measure its absorbance.
  8. Rinse the cuvette and all of the glassware you used with distilled water and return them to where you found them.
Waste Disposal. All materials can be washed down the sink with plenty of water to neutralize the acids and bases.

Calculations:

  1. Calculate the number of moles of pure acetylsalicylic acid used in your "Stock Solution"
  2. Calculate the molarity of your "Stock Solution"
  3. Calculate the molarity of each of your standard solutions, "A", "B", "C", "D", and "E".
  4. Plot the concentration of each standard solution vs. its measured absorbance (Beer's Law plot).  Use Mr. Plot (on the Saunders Interactive Chemistry CD) to provide the correct labels for the graph  (Labels for Title, Experiment, Date, X-axis, Y-axis).  Under Mr. Plot's 'Plot Options' pick the 'Least Squares Fit' option.  Mr. Plot will then draw the best straight line through your data.  It will also print the equation of that line at the bottom of the plot.
  5. Calculate the purity of your aspirin.  When you move your mouse over the plot, Mr. Plot reports the x,y position.  Move your mouse to the least squares line and move along the line until the the 'y' value is the same as the absorbance you measured for "My Aspirin" sample.  The 'x' value, is the concentration of "My Aspirin".  Divide this value by the value calculated assuming your sample was pure, and multiply by 100.
  6. If you used Excel to produce your Beers' Law plot, make sure to add a 'Trend Line'.  This 'Trend Line' is the least squares line through your data.  You will also want to set the plot options to show the equation of the line on the graph.  You can use this equation to calculate the concentration of your aspirin sample by using your absorbance value for 'y' and solving for 'x'.  Divide this 'x' value by the value calculated assuming your sample was pure, and multiply by 100.

(Updated 6/7/07 by C.R. Snelling)