Materials Included In Kit

Additional Materials Required

Prelab Preparation

1. Make ethyl acetate-hexane, 50% v/v. Mix 150 mL of ethyl acetate with 150 mL of n-hexane. Stir to dissolve.
2. Crush the aspirin tablets prior to student use.
3. Carefully cut the TLC plates to 3 x 8 cm strips.

Safety Precautions

Concentrated sulfuric acid is severely corrosive to eyes, skin and body tissue. Clean up all spills immediately. Acetic anhydride is a corrosive liquid and the vapors are highly irritating. The liquid is flammable and a strong lachrymator—contact with the liquid will cause severe eye irritation. Work with acetic anhydride in the hood or a well-ventilated lab only. Do not inhale the vapors. Ethyl alcohol and ethyl acetate are flammable liquids and severe fire risks. Do not use near flames, sparks or other ignition sources. Salicylic acid may be harmful if swallowed. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Remind students that all chemicals prepared in the lab are for laboratory use only and should never be removed from the lab. The aspirin prepared in this lab may be impure and contaminated with chemicals that could be dangerous if ingested. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The dry solids may be placed in the trash according to Flinn Suggested Disposal Method #26a. Any solutions remaining from the chemical tests of the products should be collected in a “flammable organic waste” container for licensed hazardous waste disposal.

Lab Hints

• This experiment may be completed within a typical 2-hour lab period.
• Concentrated sulfuric acid and acetic anhydride are corrosive liquids. We recommend setting out and dispensing these chemicals in a central and supervised location in the fume hood. Place the reagent bottles on a demonstration tray or laboratory spill mat to contain any chemical spills. To further reduce spillage, tape a small test tube to the side of each reagent bottle. Store a pipet in the test tube for students to use.
• TLC sheets contain a fluorescent powder that glows bright green when placed under a shortwave UV lamp. The spots that contain separated compounds will not glow green and will frequently be dark. Carefully mark the location of the spots using a pencil while the TLC plate is under the UV lamp. Trace an outline of the spots on the plate so a record of the TLC can be kept.
• Visualization of the compounds on TLC plates requires a short-wavelength UV light source—a black light is not high enough energy.
• Good technique is required for good resolution (separation) by TLC. Poor separation may be caused by: too much sample applied to the TLC plate, initial spot too large, initial spot below the solvent level in the developing jar.
• Thiele-Dennis tubes are designed to give excellent convection and heat transfer for melting point and boiling point determinations. The unique design creates convection currents when the oil inside the tube is heated, allowing the oil to flow continuously through the tube without stirring or shaking. The recommended heating fluid is silicone oil. Fill the tube to the level shown—the oil will expand when heated. Never heat a closed system! Make sure the cork or stopper in the Thiele-Dennis tube is notched to allow air to escape (see Figure 6).
  {14047_Hints_Figure_6_Design and use of a Thiele-Dennis tube}
• Melting points and boiling points of organic compounds may be found in the Merck Index, CRC Handbook of Chemistry and Physics and Lange’s Handbook of Chemistry. Melting points may also be found online in Section 9 of the Safety Data Sheets. Visit the Flinn website at www.flinnsci.com to download current SDS for all Flinn chemicals.
• The product may be recrystallized to remove impurities responsible for a slightly lower melting point compared to pure acetylsalicylic acid. Measure and record the combined mass of the watch glass, filter paper and aspirin product.
• Compare the acidity of salicylic acid and acetylsalicylic acid (aspirin) by testing the pH using narrow range (3.0–5.5) pH paper. Salicylic acid is more acidic than aspirin.
• The purity of aspirin may be analyzed by microscale acid–base titration. See the Supplementary Information in the Further Extensions section.
• (Optional) To recrystallize the product, dissolve in 3 mL of ethyl alcohol and gently heat (do not boil) the mixture. Add about 6 mL of distilled water to the hot solution until the solution is slightly cloudy. Cool the flask in an ice bath to obtain crystals and then filter and wash the recrystallized product.

Further Extensions

Supplementary Information: Microscale Titration of Aspirin
Aspirin is acetylsalicylic acid, a monoprotic weak acid. When aspirin is dissolved in water, it ionizes to give its conjugate base and H₃O⁺ ions:

A regular-strength aspirin tablet contains 325 mg of acetylsalicylic acid per tablet. This amount of acetylsalicylic acid is easily analyzed using microscale titration to determine the composition and purity.

Acetylsalicylic acid is not very soluble in water. To prepare a solution for analysis, grind one regular-strength aspirin tablet in a mortar and pestle, and dissolve the powder in about 10 mL of ethyl alcohol. Add enough distilled water to the aspirin/alcohol mixture to make 500 mL of solution. There will be some insoluble residue (starch and binder) present. This residue should not interfere with the titration.

Sample Titration Procedure

1. Place a 24-well reaction plate on top of a piece of white paper.
2. Using a microtip pipet, add 25 drops of aspirin solution to well A1, followed by 1 drop of 0.05% phenolphthalein indicator solution. For best results, do NOT add more than one drop of indicator.
3. Titrate the aspirin solution in well A1 dropwise with 0.010 M sodium hydroxide solution from a microtip pipet.
4. Count the number of drops of sodium hydroxide required to reach a pink or red-violet endpoint.
5. Repeat steps 2–4 five more times and record the number of drops of sodium hydroxide required for each trial.
6. Calculate the average number of drops of sodium hydroxide for trials 1–6.

Data Table and Calculations

Note: The last factor (0.5 L) in the above equation represents the fact that the original aspirin tablet was dissolved in 500 mL (not 1 L) of solution.

Purity of aspirin tablet = (320 mg/325 mg) x 100% = 98%

Answers to Prelab Questions

1. Acetic anhydride is a lachrymator. What is a lachrymator? What safety precautions should be followed when working with acetic anhydride?

   A lachrymator is a substance that causes tearing and severe eye irritation. Avoid breathing the vapor, and use the substance in the hood or in a well-ventilated lab only. Wear chemical splash goggles and chemical-resistant gloves and aprons.

2. Why is concentrated sulfuric acid used in this experiment? What are the hazards of working with concentrated sulfuric acid?

   Concentrated sulfuric acid is used as an acid catalyst for the esterification reaction of salicylic acid with acetic anhydride. It is extremely corrosive and will cause severe burns. Be careful when working with concentrated acids and avoid contact with eyes and skin. Clean up acid spills immediately. Wear chemical splash goggles and chemical-resistant gloves and aprons.

3. Calculate (a) the molar mass of salicylic acid and acetic anhydride and (b) the number of moles of each that will be used in this experiment. Note: The density of acetic anhydride is 1.08 g/mL.
   {14047_PreLabAnswers_Equation_2}
4. Define the term limiting reactant and identify the limiting reactant in this experiment.

   The limiting reactant in a chemical reaction is present in the smallest number of moles based on the stoichiometric mole ratios for the reactants. The number of moles of the limiting reactant determines the maximum amount of product that can be obtained. Salicylic acid is the limiting reactant in the synthesis of aspirin.

5. Determine the chemical formula of acetylsalicylic acid (aspirin) and calculate its molar mass.

   Acetylsalicylic acid, C₉H₈O₄, 180 g/mole

6. Figure 2 is a sample chromatogram for three dyes. (a) Calculate the R_f value for the spot in dye B using A as an example. (b) Dye C gave two spots on the chromatogram. What does this tell you about the composition of the sample?
   {14047_PreLabAnswers_Figure_2}

Sample Data

Laboratory Report

Answers to Questions

Laboratory Report

1. Calculate the number of moles of salicylic acid used in this experiment.
   {14047_Answers_Equation_3}
2. Calculate the maximum amount of acetylsalicylic acid in grams that may be obtained from this amount of salicylic acid. This is the theoretical yield. Hint: See Prelaboratory Questions 3–5.

   0.0036 mole x 180 g/mole = 0.65 g aspirin

3. Determine the mass of aspirin obtained in this experiment and calculate the percent yield.
   {14047_Answers_Equation_4}

   (0.30 g/0.65 g ) x 100 = 46%

4. Iron(III) ions are used as a qualitative test for phenols (aromatic compounds containing an –OH functional group). (a) What compound was used as a positive control for the Fe³⁺ test in this experiment? (b) Did the reaction product give a positive or negative test result with Fe³⁺ ions? Explain.
   1. Salicylic acid, which contains a phenolic –OH group in its structure, served as a control sample that would give a positive test. Salicylic acid gave a dark purple product when iron(III) chloride was added.
   2. The aspirin product gave a negative test with Fe³⁺ ions—no color change was observed. Aspirin does not contain an –OH functional group attached to the aromatic ring since the –OH group in salicylic acid was converted to an acetyl (–OCOCH₃) group.
5. Old aspirin tablets often have a faint vinegar (acetic acid) smell and give a positive test with iron(III) ions. Write a balanced chemical equation for the hydrolysis of aspirin (reaction of aspirin with water) to explain these observations.
   {14047_Answers_Figure_7}
   Hydrolysis of aspirin gives salicylic acid (positive Fe³⁺ test) and generates acetic acid (vinegar smell) as a byproduct.
6. Acetic anhydride was used in excess in this experiment. What does this mean, and how was the excess acetic anhydride decomposed at the end of the reaction?

   The number of moles of acetic anhydride was greater than the number of moles that would react completely with the amount of salicylic acid in the reaction mixture. Excess acetic anhydride was therefore left over in the reaction mixture after all the salicylic acid had been converted to product. The acetic anhydride remaining at the end of the reaction was decomposed by adding water to the reaction mixture prior to crystallization.

7. Look up the melting points of salicylic acid and aspirin (acetylsalicylic acid) in a reference book or online and compare with the melting point of the reaction product.

   Literature melting points:
   Salicylic acid, mp 157–159 °C
   Acetylsalicylic acid, mp 135 °C (dec)
   Melting point of product, 120–122 °C
   The product does not appear to be pure. It may be a little wet. The melting point of a mixture is lower than that of a pure substance.

8. Describe the results of TLC analysis of the aspirin product obtained in this experiment. Compare the purity and R_f values of the product against commercial aspirin and the starting material.

   The product showed only one spot on the TLC plate, with an R_f value equal to that of aspirin. Salicylic acid had a much greater Rf value than the product or commercial aspirin sample, indicating that the product was pure and did not contain any residual starting material.

Introduction

Aspirin, first synthesized in 1897, is one of the oldest yet most common drugs in use today. Like many modern drugs, aspirin has its roots in an ancient folk remedy—the use of willow extracts to treat fever and pain. Aspirin is prepared the same way today as it was more than 100 years ago. Let’s look at the structure, synthesis and properties of aspirin.

Concepts

• History of aspirin
• Salicylic acid derivatives
• Esters and esterification
• Excess and limiting reagents
• Percent yield
• Thin-layer chromatography

Background

Native Americans, as well as the ancient Chinese, Egyptians and Greeks, used willow extracts to treat fever, pain and inflammation. The Ebers papyrus, dating to at least 1500 BCE in Egypt, contains the earliest written reference to the use of willow extracts, “to draw the heat out” from inflammation. Willow extracts remained a popular folk medicine remedy throughout the Middle Ages. The first scientific study of the effectiveness of willow extracts was carried out in 1763 by the Rev. Edward Stone in England. In one of the first ever “clinical trials” of a drug, Stone reported using willow extracts to treat fever and pain in more than 50 patients suffering from malaria.

In the early 19th century, organic chemistry was only a fledgling science, with roots in the study of natural products. In 1828, Johann Büchner at the University of Munich in Germany isolated a crystalline compound from willow bark and named it salicin, after the Latin name for the white willow, Salix alba. Ten years later, the Italian chemist Raffaele Piria converted salicin to salicylic acid, which had also recently been isolated from meadowsweet (Spiraea). Salicylic acid (see Figure 1) was found to be the active ingredient responsible for the medicinal properties of many plants, including willow, poplar, aspen and myrtle. In 1859, Hermann Kolbe at Marburg University in Germany determined the chemical structure of salicylic acid and synthesized it from phenol, a derivative of coal tar. By 1870, salicylic acid was widely used in Europe for the treatment of arthritis, pain and fever. Unfortunately, the compound was “tough to swallow” and very irritating to the stomach. Many people could not tolerate the drug because of its severe and unpleasant side effects.

Felix Hoffmann, an organic chemist working at Friedrich Bayer and Company in Germany, attempted to chemically modify salicylic acid and thus reduce its side effects. In 1897, Hoffmann synthesized acetylsalicylic acid by reacting salicylic acid with acetic anhydride in the presence of an acid catalyst (Equation 1). The synthesis of acetylsalicylic acid is an example of an esterification reaction. Replacing the phenolic –OH group in salicylic acid with an acetyl or ester functional group (–OCOCH₃), makes the compound less corrosive. Acetylsalicylic acid is an effective analgesic (pain reliever) and antipyretic (fever reducer), and it is less acidic than salicylic acid. In 1899, the Bayer Company marketed acetylsalicylic acid under the trade name aspirin, with a– denoting the acetyl group and –spirin referring to Spiraea, the plant from which salicylic acid was first isolated. It is estimated that approximately 50 billion aspirin tablets are consumed per year all over the world, and that as many as one trillion (1 x 10¹²) aspirin tablets have been produced in the more than 100 years since its discovery!

The most common uses of aspirin today are for the prevention of heart attack and stroke and to relieve the pain and reduce the inflammation of arthritis. The American Heart Association recommends “an aspirin a day” to prevent a second heart attack in individuals who have had a previous heart attack or stroke. The myriad physiological effects of aspirin were explained in 1972 by Sir John Vane (Nobel Prize in Medicine, 1982). Aspirin inhibits an enzyme involved in the synthesis of prostaglandins, hormone-like “chemical messengers” that play a key role in a variety of physiological processes, including inflammation, blood clotting, labor and childbirth, and blood pressure.

Experiment Overview

The purpose of this experiment is to prepare acetylsalicylic acid (aspirin) and analyze its purity. Chemical reaction of the product with iron(III) nitrate will be used to determine if any starting material remains or if the product decomposes. The identity of the product will also be confirmed by melting point and thin-layer chromatography (TLC).

Thin-layer chromatography or TLC is a valuable technique for analyzing organic compounds. A sample is spotted onto a thin layer of silica gel or other absorbent material that has been coated onto a plastic sheet or glass plate. As solvent is allowed to seep through the absorbent via capillary action, different compounds will move up the plate at different rates, depending on their relative attraction for the absorbent versus the solvent. Partitioning of compounds gives rise to separate bands on the TLC plate. The silica gel is typically doped with a fluorescent indicator, which allows colorless organic compounds to be visualized using ultraviolet light. The distance a compound moves along the TLC plate is compared to the overall distance the solvent travels—the ratio is called the R_f value.

Materials

Prelab Questions

1. Acetic anhydride is a lachrymator. What is a lachrymator? What safety precautions should be followed when working with acetic anhydride?
2. Why is concentrated sulfuric acid used in this experiment? What are the hazards of working with concentrated sulfuric acid?
3. Calculate (a) the molar mass of salicylic acid and acetic anhydride and (b) the number of moles of each that will be used in this experiment. Note: The density of acetic anhydride is 1.08 g/mL.
4. Define the term limiting reactant and identify the limiting reactant in this experiment.
5. Determine the chemical formula of acetylsalicylic acid (aspirin) and calculate its molar mass.
6. Figure 2 is a sample chromatogram for three dyes. (a) Calculate the R_f value for the spot in dye B using A as an example. (b) Dye C gave two spots on the chromatogram. What does this tell you about the composition of the sample?
   {14047_PreLab_Figure_2}

Safety Precautions

Concentrated sulfuric acid is severely corrosive to eyes, skin and body tissue. Acetic anhydride is a corrosive liquid and the vapors are highly irritating. The liquid is flammable and a strong lachrymator—contact will cause severe eye irritation. Dispense and work with acetic anhydride in the hood and avoid breathing mist, vapors or spray. Ethyl alcohol and ethyl acetate are flammable liquids and severe fire risks. Do not use near flames, sparks or other ignition sources. Salicylic acid may be harmful if swallowed. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Preparation of Aspirin

1. Prepare a hot water bath for use in step 5 and add a boiling stone (see Figure 3). Heat the water to about 80 °C.
   {14047_Preparation_Figure_3}
2. Tare (zero) a clean 50-mL Erlenmeyer flask and add about 0.5 g of salicylic acid. Measure and record the precise mass of salicylic acid to 0.01 g.
3. Working in the hood, add 1 mL of acetic anhydride to the Erlenmeyer flask using a graduated Beral-type pipet.
4. Carefully add 2 drops of concentrated sulfuric acid to the flask.
5. Place the flask in a clamp and attach the clamp to the ring stand. Carefully lower the Erlenmeyer flask into the hot water bath. Heat the reaction mixture for 10 minutes.
6. Half-fill a 250-mL beaker with crushed ice and water for use as an ice bath. Obtain about 15 mL of distilled water in a small beaker or test tube and cool the water in the ice bath.
7. After 10 minutes, carefully lift the Erlenmeyer flask out of the hot water bath and allow to cool.
8. Add 3 mL of ice-cold water dropwise to the reaction mixture. Note: Water may react vigorously with acetic anhydride to form acetic acid. Avoid breathing the vapor.
9. Add an additional 5–6 mL of ice-cold water to the flask, and place the flask in the ice bath to crystallize the aspirin. If no crystals have formed after 5 minutes, gently scratch the sides of the flask with a stirring rod.
10. Keep the flask in the ice bath for 10 minutes to complete crystal formation.
11. Measure and record the mass of filter paper and set up a vacuum filtration apparatus as shown in Figure 4. The second filter flask is used as a trap to prevent back up of water from the aspirator to the flask when the vacuum is released.
    {14047_Procedure_Figure_4}
12. Isolate the precipitate by vacuum filtration. Careful transfer techniques are essential for accurate results.
13. Rinse any remaining crystals from the flask into the Büchner funnel with a small amount (no more than 4 mL) of ice-cold water.
14. Measure and record the mass of a clean and dry watch glass. Dry the aspirin under vacuum. Then carefully remove the filter paper and crystals and place them on the watch glass. Allow the crystals to air dry for 10 minutes.
15. Measure and record the combined mass of the watch glass, filter paper and aspirin product.

Properties of Aspirin

1. Label three small test tubes A–C, and add a small amount (about 20 mg) of salicylic acid (A), the reaction product (B) and crushed aspirin (C) to the appropriate test tube. Add about 2 mL of 50% ethyl alcohol to each test tube to dissolve the solids.
2. Add 3 drops of 0.1 M iron(III) chloride solution to each test tube and record observations.
3. Measure the melting point of the reaction product and compare against the literature values for both pure acetylsalicylic acid as well as the starting material, salicylic acid.
4. Compare the purity of the reaction product and aspirin using TLC analysis as described below (steps 5–15).
5. Dissolve about 50 mg of salicylic acid (A), the reaction product (B) and crushed aspirin (C) in 10 mL of ethyl alcohol in separate, appropriately labeled medium test tubes. Cap each test tube with a cork or rubber stopper and swirl to dissolve the solids.
6. Obtain or prepare a 3 x 8 cm TLC plate and mark it with a pencil as shown in Figure 5.
   {14047_Procedure_Figure_5_Thin layer chromatography}
7. Add about 25 mL of a 50:50 ethyl acetate–hexane solution to a chromatography jar or beaker. The liquid level should cover the bottom of the container to a depth of 0.5 cm. Do not use too much solvent! Cover the container and allow it to stand for 15 minutes to equilibrate the chamber with solvent vapor.
8. Spot samples A−C onto the TLC plate. Using a microcapillary tube, place one small drop of sample solution A on the first dot. Let the spot dry completely, then place another drop of the same sample on top of the first drop, and again let it dry. Repeat until 5−6 drops of the sample have been applied to the same spot. Allow to dry completely.
9. Repeat step 8 to spot samples B and C on the TLC plate.
10. Pick up the TLC plate by the edges, and carefully place it in the chromatography chamber. The sample spots on the TLC plate must be above the level of liquid in the chromatography chamber. Cover the chamber without disturbing the plates.
11. Do not disturb or move the jar while the chromatogram is developing!
12. The solvent (liquid) will be drawn up the TLC plate by capillary action. When the liquid level has moved to within 1 cm of the top of the plate, remove and place it on a paper towel. Use a pencil to mark the position of the solvent front on the plate.
13. Allow the plate to air dry.
14. View the plate using a shortwave UV light to visualize the spots. Outline each spot using a pencil, and mark its center.
15. Measure the distance from the starting line to the center of each spot. Record the distances and calculate the R_f value for each sample.

Student Worksheet PDF

14047_Student1.pdf