ANALYSIS OF ASPIRIN: SPECTROPHOTOMETRY
The purpose of this experiment is to evaluate the percent aspirin on a commercial aspirin tablet using an
instrumental method, spectrophotometry.
In a spectrophotometer, light from a strong lamp passes through a monochromator, which breaks the light
into its component colors using a grating, then uses mirrors and slits to direct a light beam of the desired
wavelength through the sample compartment, where you place a tube (cuvette) of solution. A detector on
the other side of the sample compartment measures the amount of light that passes through the solution
and compares that to the amount of light that passes through a control solution.
The result is displayed as either transmittance or absorbance. Transmittance is the ratio of light passing
through the sample (I) to light passing through the control (I0). Absorbance is the negative logarithm of
transmittance:
I
I
T
A   logT   log
I0
I0
The absorbance of a solution depends on three things: the nature of the solution, the distance the light
travels through the solution (called the path length), and the concentration of the solution. This is
summarized by the Beer-Lambert Law, A = abc
• A is the measured absorbance
• a is a constant (the absorptivity); it depends on  and the nature of the solution
• b is the path length
• c is the solution concentration
If the type of material and the cuvette diameter are always the same, ab is constant and the measured
absorbance depends directly on the concentration of the solution. The absorbance is most sensitive to
concentration if measured at a wavelength of high absorbance for that material.
To use the spectrophotometer, turn it on and let it warm up for 5 minutes, so the lamp and detector
outputs are stable. Select the desired wavelength, and use one knob (usually on the left) to set the detector
to read A = ∞ or T = 0% when no light passes through the sample compartment (this is called the dark
current). Put a cuvette of control solution in the sample compartment, and adjust the other knob (usually
on the right) to read A = 0 or T = 100%. The detector is now set for the range of light possible for your
experiment. Put a cuvette of sample solution into the sample compartment and record the absorbance or
transmittance of the sample.
To prepare the aspirin, you will hydrolyze it with NaOH, then react it with acidic iron (III) ions. The Fe3+
ions and salicylate ions form a deep violet complex, which absorbs strongly at 530 nm.
Procedure
Prepare iron (III)–salicylate complex solutions of known concentration (your standards)
1. Weigh about 0.16 g pure acetylsalicylic acid into a clean 50 mL flask and record the exact mass used.
Add 5 mL of 1 M NaOH solution and heat to boiling, being careful to avoid spattering.
Allow the reaction mixture to cool, then transfer all of it into a 100 mL volumetric flask, using
deionized water from your wash bottle to assure complete transfer. Add deionized water to bring the
solution to exactly 100 mL. Mix thoroughly. This is your standard stock solution.
2. Label 5 small clean dry flasks A – E. Use a 1–mL pipet to add 0.50 mL of the standard stock solution
into flask A, then use your 10–mL Mohr pipet to add 9.50 mL acidified 0.02 M FeCl3 solution, for a
total of exactly 10.00 mL. Mix thoroughly. Note the exact solution volumes used.
3. Prepare standards B – E as in step 2, using 0.40 mL, 0.30 mL, 0.20 mL, and 0.10 mL of standard stock
solution and an appropriate amount of acidified 0.02 M FeCl3 solution to give a total volume of 10.00
mL. If you goof while pipetting and add a bit too much or too little, you don't need to repeat that
sample: just record the exact volumes used.
Prepare your aspirin sample
4. Perform the same hydrolysis as in step 1, using about 0.16 g of ground aspirin tablet. Make one
dilution of the sample stock solution, using 0.30 mL of your sample stock solution + 9.70 mL
acidified 0.02 M FeCl3 solution.
Determine of absorbance of each solution
5. Set up the spectrophotometer at 530 nm, as described in the introduction. Use acidified 0.02 M FeCl3
solution (the diluant) as the control solution. Read the absorbance of each standard A – E, and of the
diluted aspirin sample (if your spectrophotometer has an analog readout, read %T and convert later).
You need not rinse the cuvette between samples (in fact, leftover water introduces more error than the
same amount of leftover solution). Instead, use a disposable Beral pipet to withdraw as much solution
as possible, then go on to the next solution.
Analysis and Discussion
Calculate the concentration of each standard, in grams acetylsalicylic acid per mL. Convert %T to
absorbance, if necessary, then plot the absorbance of standards A – E as a function of concentration.
Determine the slope and correlation coefficient of the graph. Mark the absorbance of the diluted sample
on the standard curve and determine its concentration from the slope. Calculate the concentration of
aspirin in the sample stock solution, and compare that to the mass used. Calculate the % by mass of
acetylsalicylic acid in your sample. As always, be attentive to significant figures.
If you used a commercial aspirin tablet, calculate the % acetylsalicylic acid in an aspirin tablet from
information given on the aspirin bottle (you may need to weigh a few aspirin tablets), and compare this
calculated value to your results.
Discuss the precision of your standard curve, including both the agreement between the points and the
line and the value of the y intercept. Discuss the accuracy of your result.