1.13: Experiment_613_Spectrophotometric Determination of Aspirin_1_2_2

Student Name		Laboratory Date: Date Report Submitted:	___________________________
Student ID		Experiment Number and Title	Experiment 613: Spectrophotometric Determination of Aspirin

Experiment 613 Spectrophotometric Determination of Aspirin

Section 1: Purpose and Summary

Most aspirin tablets are said to contain 5 grains of the active ingredient acetylsalicylic acid. (The grain is a unit of mass and is equal to 0.0647989 grams.) How reliable is this figure? This experiment will enable you to check up on the manufacturer.

The method of analysis that we will use is called "spectrometry" or "spectrophotometry". It depends on the fact that molecules can absorb electromagnetic radiation of certain wavelengths, using the energy of the radiation to "excite" electrons in their atoms. (Thus they have absorption spectra, just as isolated gas atoms do.) The greater the concentration of a particular molecule present in a sample, the more light of a particular wavelength the sample will be able to absorb. The absorbance of light by the

sample increases in direct proportion to the concentration of the molecule present. We will use electromagnetic radiation ("light") of wavelength (λ) 297 nm, which falls in the ultraviolet range. In this experiment, 297 nm light shines through the sample, and the amount of light absorbed by the sample is measured.

The chemical name of the active ingredient in aspirin is acetylsalicylic acid. In order to dissolve aspirin completely in water, you will convert it chemically to the salt sodium salicylate before measuring the absorbance. Notice that 1.00 mole of sodium salicylate is produced for every 1.00 mole of

acetylsalicylic acid (aspirin) used up:

Sodium salicylate will then be reacted with acidic Fe³⁺ to form salicylatoiron(III) complex, [FeSal]⁺. This complex displays a maximum absorption at a wavelength of 525 nm and has a purplish red color. The absorption of the [FeSal]⁺ complex will be directly proportional to the concentration of salicylate in the sample, but it will not tell us the actual concentration in any given sample. In order to be able to convert an instrument reading to an actual concentration of salicylate, we must first calibrate the instrument using solutions of known sodium salicylate concentration.

Overall, this experimental project consists of:

• Carefully preparing 5 solutions of known concentrations
• Measuring the absorbance of each known and unknown solution

• Carefully preparing an aspirin solution (of unknown concentration)
• (After lab) Preparing a calibration curve from the data of the 5 solutions of known concentrations
• (After lab) Using the calibration curve to calculate the amount of aspirin in one tablet or pill.

Section 2: Safety Precautions and Waste Disposal

Safety Precautions:

All of the materials in this experiment are relatively harmless. Aspirin obtained in lab should not be ingested. Sodium hydroxide solutions are corrosive. Use of eye protection is recommended for all experimental procedures.

Waste Disposal:

Dispose of solutions down the sink drain with plenty of water. Dispose in the trash any filter paper and undissolved residue from the aspirin tablet.

Section 3: Procedure

Part 1: Preparing solutions of known concentrations

Ask your instructor to describe and demonstrate the use of volumetric pipets, volumetric flasks, and burets. Into your smallest beaker, obtain 10-15 mL of the standard solution of approximately 0.03 M sodium salicylate. Record the actual concentration (include all the significant figures).	Concentration (Molarity) of the Standard Solution of sodium salicylate:
Using a rubber bulb and a 5.00 mL volumetric pipet, carefully measure 5.00 mL of the standard solution and place it in a 100.0 mL volumetric flask. Carefully fill the flask with laboratory water exactly to the 100 mL mark. Mix the sample thoroughly by carefully inverting the stoppered flask repeatedly for at least 2 minutes. (Hold the glass or rubber stopper tightly in place with your thumb. Make sure no liquid leaks out before the solution is completely mixed.) This is your Stock Solution.
Obtain 6 clean and dry beakers that are larger than 100 mL. Label these beakers “Stock”, “1”, “2”, “3”, “4”, and “5”.
Empty the 100.0 mL volumetric flask into the beaker labeled “Stock”. Rinse the volumetric flask with plenty of laboratory water.
Fill a 50 mL buret using your Stock Solution. Fill it close to the 0.00 mL mark, making sure that you can accurately read the volume at the meniscus (the lowest part of the surface of the solution inside the buret). It is OK if the volume is not exactly 0.00 mL.	Initial Buret Reading (mL): (Include 2 decimal places!)
Now prepare the first known solution, Solution #1. Place the 100 mL volumetric flask below the buret, then turn the stopcock of the buret so that about 5 mL drains into the volumetric flask. Afterwards, read the buret.	Buret Reading #1 (mL): (Include 2 decimal places!)
Quick check: Subtract the Initial Buret Reading from Buret Reading #1. This number should be close to 5 mL. Later you will use this volume to calculate the concentration (molarity) of sodium salicylate in Solution #1. Let’s name this volume “Transfer Volume #1”	Calculate the difference between the Initial Buret Reading and Buret Reading #1. (Include 2 decimal places!) “Transfer Volume #1”: ______________ mL
To the 5 mL contained in the volumetric flask, carefully fill the flask with 0.010 M FeCl3 in 0.1 M HCl solution the 100 mL mark and mix well. This is Solution #1. Empty the 100 mL volumetric flask into the beaker labeled “1”. Rinse the volumetric flask with plenty of laboratory water.
Now prepare the second known solution, Solution #2. Place the 100 mL volumetric flask below the buret, then turn the stopcock of the buret so that about 10 mL drains into the volumetric flask. Afterwards, read the buret.	Buret Reading #2 (mL): (Include 2 decimal places!)
Subtract Buret Reading #1 from Buret Reading #2. This number should be close to 10 mL. Record Transfer Volume #2. Carefully fill the volumetric flask with 0.010 M FeCl3 in 0.1 M HCl solution exactly to the 100 mL mark and mix well. This is Solution #2. Empty the 100 mL volumetric flask into the beaker labeled “2”. Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between Buret Readings #1 and #2. (Include 2 decimal places!) “Transfer Volume #2”: ______________ mL
Now prepare the third known solution, Solution #3. Place the 100 mL volumetric flask below the buret, then turn the stopcock of the buret so that about 15 mL drains into the volumetric flask. Afterwards, read the buret.	Buret Reading #3 (mL): (Include 2 decimal places!)
Subtract Buret Reading #2 from Buret Reading #3. This number should be close to 15 mL. Record Transfer Volume #3. Carefully fill the volumetric flask with 0.010 M FeCl3 in 0.1 M HCl solution exactly to the 100 mL mark and mix well. This is Solution #3. Empty the 100 mL volumetric flask into the beaker labeled “3”. Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between Buret Readings #2 and #3. (Include 2 decimal places!) “Transfer Volume #3”: ______________ mL
Refill the buret with your stock solution (use the beaker labeled “Stock”). Read the buret and record this volume (it should be close to the 0.00 mL mark). It is OK if the volume is not exactly 0.00 mL.	Initial Buret Reading After REFILL (mL): (Include 2 decimal places!)
Now prepare the fourth known solution, Solution #4. Place the 100 mL volumetric flask below the buret, then turn the stopcock of the buret so that about 20 mL drains into the volumetric flask. Afterwards, read the buret.	Buret Reading #4 (mL): (Include 2 decimal places!)
Subtract the Initial Buret Reading After REFILL from Buret Reading #4. This number should be close to 20 mL. Record Transfer Volume #4. Carefully fill the volumetric flask with 0.010 M FeCl3 in 0.1 M HCl solution exactly to the 100 mL mark and mix well. This is Solution #4. Empty the 100 mL volumetric flask into the beaker labeled “4”. Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between Buret Reading #4 and the Initial Buret Reading After REFILL. (Include 2 decimal places!) “Transfer Volume #4”: ______________ mL
Finally, prepare the fifth known solution, Solution #5. Place the 100 mL volumetric flask below the buret, then turn the stopcock of the buret so that about 25 mL drains into the volumetric flask. Afterwards, read the buret.	Buret Reading #5 (mL): (Include 2 decimal places!)
Subtract Buret Reading #5 from Buret Reading #4. This number should be close to 25 mL. Record Transfer Volume #5. Carefully fill the volumetric flask with 0.010 M FeCl3 in 0.1 M HCl solution exactly to the 100 mL mark and mix well. This is Solution #5. Empty the 100 mL volumetric flask into the beaker labeled “5”. Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between Buret Readings #4 and #5. (Include 2 decimal places!) “Transfer Volume #5”: ______________ mL
You should now have 5 beakers labeled 1-5, each containing 100 mL of sodium salicylate at different concentrations.

Part 2: Using an aspirin tablet to prepare a solution of unknown concentration

Crush one aspirin tablet in a clean porcelain mortar. Read the label of the bottle of Aspirin to find the number of GRAINS in each tablet. If Grains are not reported, record the number of grams instead.	From the Aspirin bottle, report the number of GRAINS (if not available, then report grams) each tablet should contain:
Transfer all the solid to a beaker (any size over 100 mL). Use several portions of 0.1 M NaOH (about 50 mL total) to rinse all the solid into the beaker. Be sure to rinse both mortar and pestle.
Gravity Filtration: Prepare a folded filter paper (coffee filter paper works!) in a funnel and stand the funnel in the neck of a clean 250.0 mL volumetric flask. (The flask must be clean but does not have to be dry.) Make sure no particles of dust, paper, or other impurities get into the flask. Pour the aspirin-containing liquid from the beaker onto the filter paper in the funnel. The clear solution that comes through the filter contains your sample; the solid that is left on the filter paper will eventually be discarded. Rinse the beaker three times with 10 mL of 0.1 M NaOH (about 30 mL total) and pour these rinsings into the filter funnel also.
When the sample has been filtered, remove the funnel, and dispose in the trash the filter paper and undissolved residue from the aspirin tablet.
Carefully add enough laboratory water to bring the solution level exactly to the 250.0 mL line on the neck of the flask.
Mix the sample thoroughly by carefully inverting the stoppered flask repeatedly for at least 2 minutes. (Hold the glass or rubber stopper tightly in place with your thumb. Make sure no liquid leaks out before the solution is completely mixed.)
The solution at this point is too concentrated to be used. Prepare a dilution of it by using a volumetric pipet to transfer exactly 2.00 mL of it into a clean 100.0 mL volumetric flask (the flask may be wet as long as it is clean). Carefully add 0.010 M FeCl3 in 0.1 M HCl solution to bring the solution level up to the 100.0-mL mark on the flask. Mix thoroughly.
Label this 100 mL volumetric flask as “unknown”.

Part 3: Measuring the absorption of UV light by solutions of sodium salicylate

1. The proper use and operation of the spectrophotometer can be found in Appendix: Technique I Use of Spec 20/Spec 200.

2. Adjust the spectrophotometer to measure the absorbance at 525 nm.

3. Use 0.010 M FeCl3 in 0.1 M HCl solution as the reference liquid - that is, set the instrument for zero absorbance when the light path is passing through this iron(III) solution.

Measure the absorbance for each of your solutions of known concentrations (beakers 1-5) and for your unknown solution (the 100 mL volumetric flask labeled “unknown”). Rinse the cuvette (sample holder) thoroughly with laboratory water, then with each solution in between measurements. Record your data in the table below.

Sample	Absorbance at 525 nm
Solution #1
Solution #2
Solution #3
Solution #4
Solution #5
Unknown

Quick check:

• The absorbance should increase from Solution #1 (lowest) to Solution #5 (highest).
• The absorbance of the unknown should lie between the absorbances of Solution #1 and Solution #5.

Section 4: Calculations

(These may be completed after lab. However if time is available, it is recommended that you work on these calculations before leaving lab because if an error is found in your data, you may have time to obtain better data in lab.)

1. Calculate the concentration of your known solutions using the Dilution equation: M₁V₁ = M₂V₂.

Name of Solution	M1 = Concentration of the solution to be transferred (the initial solution)	V1 = Volume transferred	M2 = Final Concentration (Calculate this number with correct SIGNIFICANT FIGURES)	V2 = Final Volume of the solution after dilution is completed
Stock Solution	Concentration (Molarity) of the Standard Solution of sodium salicylate (from label on the standard bottle) =	5.00 mL		100.0 mL
Solution #1	Concentration of your Stock Solution=	Transfer Volume #1 =		100.0 mL
Solution #2	Concentration of your Stock Solution=	Transfer Volume #2 =		100.0 mL
Solution #3	Concentration of your Stock Solution=	Transfer Volume #3 =		100.0 mL
Solution #4	Concentration of your Stock Solution=	Transfer Volume #4 =		100.0 mL
Solution #5	Concentration of your Stock Solution=	Transfer Volume #5 =		100.0 mL

2. Construct the Calibration Curve

Using Microsoft Excel or another spreadsheet software package, plot the Final Concentration (y-values) vs. absorbance (x-values) for Solutions 1-5. Use Microsoft Excel to find the Line of Best Fit (this is also called Linear Regression).

Basic Instructions for Graphing with Microsoft Excel^®

1. Input data into spreadsheet with x-coordinate data in the first column and y-coordinate data in the second column.

2. Select data columns with mouse.

3. Select chart wizard from either the toolbar or under the insert menu.

4. Select desired graph type. Usually xy scatter graph.

5. Select next button.

6. Select data range if required (usually not needed).

7. Select next button.

8. Input graph title and data labels.

9. Select next button.

10. Choose insert graph as new sheet.

11. Select data points with mouse.

12. Add trend line under chart menu or right click.

13. Select trend line type (linear).

14. Use the options tab to include the linear equation and r² value on the graph.

15. Make the graph look pretty using format plot area.

16. Double click axes to adjust limits.

Write the equation of the line obtained from Linear Regression (step 14 above):

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Ask your instructor if a plot of your data should be included with your lab report.

3. Using your equation written above, calculate the concentration of the Unknown Solution:

Sample	Absorbance at 525 nm	Concentration
Unknown

4. From the concentration of the unknown solution, calculate the amount of aspirin in the tablet.

Copy the concentration of the unknown solution (from table above) that was placed into the spectrophotometer.	Concentration of the unknown solution:
In lab, you prepared a dilution of the unknown solution by using a volumetric pipet to transfer exactly 2.00 mL of it into a clean 100.0-mL volumetric flask. Let’s calculate the concentration of that 2.00 mL solution using M1V1 = M2V2. M1 = (calculate this) V1 = 2.00 mL M2 = Concentration of the Unknown Solution (from above) V2 = 100.0 mL	Concentration of the unknown solution before dilution:
That 2.00 mL solution came from the 250.0 mL volumetric flask (this volume is equal to 0.2500 L). Calculate the number of moles of sodium salicylate in the 250.0 mL volumetric flask by multiplying 0.2500 L by the Concentration of the Unknown Solution Before Dilution	Moles of sodium salicylate:
For every mole of sodium salicylate in the original solution, how many moles of aspirin (acetylsalicylic acid) were there in the tablet? (See the balanced equation at the beginning of this experiment)	Moles of aspirin:
Calculate the number of grams from the number of moles using the molar mass of Aspirin (180.158 g/mol).	Grams of aspirin:
Convert grams to grains (15.43 grains = 1.000 gram).	Grains of aspirin:
Calculate the percent error using either grains or grams of aspirin: (Your grains of Aspirin – Bottle’s grains of Aspirin) * 100 Bottle’s grains of Aspirin	Percent Error:

Post-lab Questions:

1. Discuss your experimental results: do your results agree with the Aspirin manufacturer’s claim regarding the grains (or grams) of Aspirin in each tablet? Discuss any of your experimental errors that may affect your conclusions.

2. Discuss how the structure of aspirin becomes more water soluble after reacting with sodium hydroxide.

3. In a different experiment, known solutions of ibuprofen were used to create a calibration curve. The line of best fit for this calibration curve was

y = 0.000924x + 0.0000343

where x = absorbance and y = concentration (Molarity) of ibuprofen in solution.

1. If the absorbance = 0.207, what is the concentration of ibuprofen in solution?

2. Predict the absorbance for a 5.62 x 10^-4 M solution of ibuprofen.

3. Suppose an ibuprofen tablet is crushed, dissolved in sodium hydroxide, and filtered. This solution is diluted to a total volume of 5.00 L and mixed thoroughly. If this solution has an absorbance of 0.163, how many milligrams of ibuprofen were in the tablet? The molar mass of ibuprofen is 206.29 g/mol.