Aspirin lab Essay

The medicinal properties of acetylsalicylic acid have been known for millenia. The ancient Greeks noted that extracts of willow and poplar bark were effective painkillers, though it was not until the 1800s that its active ingredient, salicylic acid, was isolated. However, it was strongly acidic and irritated membranes of the mouth and stomach. In 1875, sodium salicylate was introduced in an attempt to weaken the salicylic acid, but it was found to be medically inefficient.

Above: The molecular structures of salicylic acid and sodium salicylate

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In 1893, Felix Hoffman Jr.

discovered a way to synthesize an ester of salicylic acid for Bayer Laboratories in Germany, acetylsalicylic acid. It was not as strong as salicylic acid but had the same medicinal properties. Acetylsalicylic acid was then patented in 1899 and became commercially available in 1915 under the name of Aspirin.

Left: The molecular structure of acetylsalicylic acid, commonly known as aspirin.

Today, Aspirin is one of the most commonly used drugs in the world. It is a powerful as a pain reliever, fever reducer, and swelling-reducing drug.

It also reduces the clotting of blood and can be used in small doses to prevent heart attack and stroke. However, in some cases, it may still cause irritation to the stomach membranes, and blood in the stomach lining is lost due to the consumption of aspirin. The drug may also interfere with normal blood clotting and cause Reye’s syndrome in children.

Above: An esterification reaction, which forms aspirin as one of its products.

In this experiment, it will be attempted to first prepare aspirin by chemical synthesis from salicylic acid, by acetylation with acetic anhydride. This synthesis is an esterification reaction between an ester and an acid to form a more complex ester. Esters are a type of acid in which the hyrdroxide groups are replaced with alkyl chains composed of carbon and hydrogen. Afterwards the product will be recrystallized to purify the product by allowing large, pure crystals to form by cooling slowly. Generally, the amount of product exceeds the amount of impurity, therefore the crystals of the product will form first and leave a greater ratio of impurity in the solution. Finally melting point and percentage yield will be determined in order to determine the purity of the aspirin.


Acetic anhydride
Salicylic acid (solid)
Sulphuric acid (Note: sulfuric acid will result in a lower yield, being a strong oxidizer. However sulfuric acid is still used sometimes for reactions that require the loss of water because it absorbs water while phosphoric acid does not.) Ice

Distilled water
Ethyl alcohol



Large beaker (~600mL)
Smaller beaker (~100mL)
Large test tube

Measuring Devices
Graduated cylinder

Other Equipment
Hot plate
Melttemp apparatus and at least four capillary tubes
Stirring rod
Watch glass
Buchner funnel
At least two pieces filter paper
Wash bottle
Dessicator with silica

Above: Dessicator with silica.

Above: Melttemp apparatus.


Synthesis of Aspirin
1. Heat a large beaker full of water on the hot plate to 75-85˚C. Leave the hot plate on and the beaker on the hot plate for step 5. 2. Meanwhile, weigh out 3.0 g of salicylic acid and deposit in a large test tube. 3. Using a graduated cylinder, measure 6.0 mL of acetic anhydride and deposit into the test tube. 4. Add 10 drops of 85% sulphuric acid to the flask and stir with a stirring rod. The acid will act as a catalyst to speed up the reaction. 5. Place test tube in the large heated beaker of water on the hot plate, stirring occasionally for about 10 minutes.

Crystalizing the Aspirin
1. Remove beaker from heat. Using the dropper, carefully add 20 drops of cold distilled water to the test tube. 2. Allow the mixture to cool to room temperature so that crystals may form slowly. 3. Cool the mixture further
by placing the test tube in a large beaker full of ice water to crystallize and minimize solubility for around 15 minutes, or until crystallization is complete. If crystallization does not start, tap the test tube with a stirring rod. 4. Meanwhile, weigh the filter paper with a watch glass and fill a wash bottle with around 25 mL of chilled distilled water to wash the crystals. Cool water must be used because warm water will dissolve the crystals. 5. Filter the solid aspirin through filter paper with a Buchner funnel and aspirator. 6. Rinse the crystals and test tube with the chilled water. At this point, impurities will still remain in the aspirin, necessitating recrystallization afterwards. 7. Place filter with product in a watch glass to dry overnight. 8. Weigh and measure the melting point range of the crude aspirin twice(see section below on finding the melting point).

Above: Filtering the solid aspirin using a Buchner funnel and aspirator.

Above: Crude aspirin.

Recrystallizing the Crude Aspirin
1. Dissolve crude aspirin in 10 mL of 95% ethyl alcohol in a 100 mL beaker. 2. Warm the mixture in a hot water bath on a hot plate to encourage dissolving. Note: Ethanol boils at 78˚C. DO NOT LET THE MIXTURE BOIL.

3. If the mixture does not dissolve, add an additional 2 mL of ethyl alcohol. 4. When all the aspirin has dissolved, pour in 10 mL of lukewarm distilled water until solution becomes transparent. 5. Cover the beaker with a watch glass and set aside to cool slowly undisturbed overnight so large crystals may form. 6. If an “oil” appears instead of a solid, reheat the beaker in hot water until the oil disappears. Repeat step 5. If crystals still do not appear, scratch the bottom of the beaker with a stirring rod to induce crystallization. 7. Collect crystals using vacuum filtering.

8. Rinse collected crystals with cold distilled water.
9. Allow the crystals to dry using vacuum filtering. If necessary, use a dessicator to dry out the aspirin. 10. At this time it is appropriate to
weigh the purer aspirin sample and calculate the percentage yield out of a maximum yield of 3.9g again.

Above: Dissolving crude aspirin in ethyl alcohol.

Finding the Melting Point Range
1. Place about 5 mm of aspirin crystals into a capillary tube, which is then placed into the melttemp apparatus. 2. Insert a mercury thermometer through the top.
3. Heat the apparatus until 15˚C from the expected melting point (135˚C) at which temperature the heat should be reduced. 4. Record the range of the melting point as the temperature at which the first drop of liquid appears up to when all the aspirin has been converted into a liquid. 5. Repeat the process once more to find the melting point range.

Above: View through a melttemp apparatus. Here the aspirin crystals are solid.

Safety Precautions

Acetic Anhydride
Acetic anhydride is an irritant and also flammable, therefore gloves and goggles should be worn at all times during the experiment. It is reactive to water, so in the case of fire, alcohol foam or carbon dioxide is preferred to use as an extinguisher. This chemical has harmful fumes and use of a fume hood is strongly recommended.

Sulphuric and Salicylic Acid
These acids may irritate the skin in high concentrations. Take appropriate care to avoid contact.

Handle all hot equipments with caution and never leave the Bunsen burner flame unattended. The hot beakers and test tubes must be handled with care and should only be moved using tongs. Also, tie back long hair to prevent any accidental fire, and be familiar with the location of the fire


Mass of Aspirin Synthesized

Mass of Watch Glass and Filter Paper (g)
Mass of Watch Glass, Filter Paper, and Product (g)
Mass of Product (g)
Crude Product
Final Product

Melting Point
First Appearance of Liquid (˚C)
Completely Melted (˚C)
Crude 1
Crude 2
Crude Average
Final 1
Final 2
Final Average
Theoretical Melting Range

Above: The final product.


Percentage Yield

Maximum Yield
Moles of a molecule = (Mass of substance) / (Molar mass)
Acetic Anhydride (C4H6O3): 102.09 g/mol,
Density is 1.082 g/mL, therefore 6mL contains 6mL x 1.082g/mL = 6.492 g acetic anhydride Therefore there are 6.492 g / 102.09 g/mol = 0.06359 moles
Salicylic Acid (C7H6O3): 138.12 g/mol, thus there are 3 g / 138.12 g/mol = 0.02172 moles

The balanced equation is
C4H6O3 + C4H6O3 → C9H8O4 + C2H4O2

which can also be written as
to show the molecular structure of aspirin more clearly.

Equal amounts of acetic anhydride and salicylic acid are required for this reaction, therefore salicylic acid is the limiting reagent. Therefore only 0.02172 moles of aspirin will be produced.

The molecular mass of aspirin is 180.16 g/mol

Moles of a molecule = (Mass of substance) / (Molar mass)
Mass of a substance = (Moles of a substance) * (Molar mass)

180.16 g/mol x 0.02172 mol = 3.913g aspirin

The percentage yield for this synthesis of aspirin may be calculated as follows:

Crude Product
Maximum yield = 3.913 g
Actual yield = 2.96 g
Percentage yield = Actual yield / Maximum yield = 75.65%

Final Product
Maximum yield = 3.913 g
Actual yield = 2.67 g
Percentage yield = Actual yield / Maximum yield = 68.23%

Melting Range Percentage Error
As the melting point is a range, the average value of the range will be found before calculating percentage error. Percentage error is calculated as |Theoretical value – Actual value| / Theoretical value.

Crude Product
Average of actual melting range = 120.25°C
Expected melting point = 135°C
Percentage error = |135 – 120.25| / 135 = 0.1093 = 10.93%

Final Product
Average of actual melting range = 133.5°C
Expected melting point = 135°C
Percentage error = |135 – 133.5| / 135 = 0.0111 = 1.11%

Conclusion and Analysis
The main purpose of this lab is to obtain aspirin through the chemical synthesis of salicylic acid by acetylation with acetic anhydride and crystallization. Crystallization is the process of arranging atoms or molecules in a liquid state into an ordered solid state. During this process, the sample that is composed of more than one substance is transformed into new samples, each of which consists of a single substance. Thus a pure sample of the compound is obtained.

In our experiment, pure aspirin was obtained after filtering out the impurities and excess reagent through filter paper. A method to check a solid compound’s purity after re-crystallization is to check its melting point. Impurities will always lower the melting point of a sample. In our experiment, the melting point range of the product was observed to be 129.5°C to 137.5°C. The percentage yield of the final product was calculated to be 68.23%.

It can be concluded that by recrystallizing the crude aspirin, though it decreased the percentage yield from 75.65% to 68.23%,a significantly purer aspirin was produced, as shown by the more feasible and realistic melting range. Aspirin has a theoretical melting range of 134-136°C. The crude product had a melting range of 117.5°C-123°C, and the final product one of 129.5°C-137.5°C. The percentage error of the crude product was 10.93%, while that of the final product was significantly lower at 1.11%.

Above: The 3D crystal structure of aspirin

By recrystalizing the crude aspirin slowly, it was possible to obtain large crystals with a rigorous structure by allowing the aspirin molecules to join together in a precise manner. The regular molecular crystal structure of the final product makes it more difficult for impurities to be included, eliminating many impurities present in the amorphous crude product.

The synthesis of aspirin demonstrated here demonstrates several relevant
suggestions that should be considered in the synthesis of a substance.

Firstly, the recrystallization of a product should be performed if possible and relevant. In most reactions, the amount of product exceeds the impurities. As a result, the crystals of the product should form first by virtue of greater volume. The remaining solution will be left with a greater concentration of impurities.

Secondly, melting point is a useful method of determining the purity of a product. When the theoretical melting range of a product is known, the melting range of a final product is usually narrower and closer to the theoretical melting range than the melting range of the crude product.

Finally for practical purposes, it is suggested that a beaker rather than a test tube be used in the synthesis of aspirin. Although a slanted test tube reduced the evaporation of reactants during the synthesis of aspirin, it was later found that the thin shape of the test tube made it difficult to seperate the product from the sides of the container using a scoopula.

Suggested Modifications to Procedure
If possible, when determining the melting point, the temperature should rise at a rate of 1-2°C per minute when nearing the expected melting point in order to acquire a more accurate range. The humidity of the lab should be kept at a minimum, as acetic anhydride, a reactant in the synthesis of aspirin, tends to react with water vapour to form acetic acid.

Phosphoric acid can be used instead of sulfuric acid if desired to obtain a higher yield, as sulfuric acid reacts more readily with the organic molecules involved in the reaction than phosphoric acid. However phosphoric acid does not absorb the water in the reaction, thus it may be a slower process.

Sources of Experimental Error
As in all experiments, experimental error is inevitably present, yielding inaccurate results. The complexity of this experiment only increases the
possibility of the entrance of an unforseenable and/or uncontrollable variable. This section aims to present the sources and implications of these variables.

During the course of the experiment, it was possible that the thermometer inaccurately measured the melting range of the aspirin, being non-digital in operation. Exact values and decimal points could not be visually obtained; however, it is believed that the mercury thermometer was fairly functional and accurate. The melttemp apparatus, however, limited accuracy and precision of the melting range made it very difficult to control the pace of heating of the sample to small increments of 1-2°C per minute.

Secondly, the digital scale was also a potential source of error, as it was extremely sensitive and results were easily disturbed by small fluctuations such as the flow of air in the lab. The scale was also neither calibrated nor standardized. The masses of the products or reactants are subject to such uncontrollable fluctuations. As the amounts of reactants and aspirin synthesized were relatively small, even minor fluctuations can translate into multiple percentages in error.

Thirdly, impurities in the reactants or the aspirin could not be isolated, controlled, or eliminated apart from qualitatively during recrystallization. The lab in which the experiment was performed nor the cupboard in which the product was dried are unlikely to be sterile, therefore impurities could have entered the reaction or reactants at any time, especially since the containers used in the experiment were largely left open. Both acetic anhydride and acetylsalicylic acid decompose in humid air, which reduced percentage yield. Another possible source of contamination was the lab equipments used in this experiment. Although all the lab equipments were previously rinsed with water and dried, it is possible that chemical residue from previous experiments were present during the course of the experiment.

Fourthly, because sulfuric acid was used instead of phosphoric acid as a catalyst to synthesize aspirin, the percentage yield was quite low. This occurs because sulfuric acid reacts more strongly with the organic molecules
in the reaction than phosphoric acid. However, it was still used because the synthesis of asprin entailed the absorbtion of water. Phosphoric acid does not absorb water, however sulfuric acid does. This property is called hygroscopy, and extends to dessicators such as calcium chloride and, to a lesser extent, silica, which was used to dry the aspirin prior to testing the melting point of the crude product.

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Experiment 11Synthesis and Analysis of Aspirin
Chemicals & Equipment
Bunsen burner or hot platesalicylic acid
Boiling stonescommercial aspirin
Buchner funnel1% FeCl3
500-mL filter flaskacetic anhydride
filter paper18M H2SO4
250-mL flask & beaker95% ethanol
melting point apparatuscapillary tubes
Time: Part I & Part II (3 hours)

Aspirin is most widely sold over-the-counter drug. It has the ability to
reduce fever (an antipyretic), to reduce pain (an analgesic), and to reduce swelling, soreness, and redness (an anti-inflammatory agent). One of the first recorded accounts for the discovery of aspirin appeared in England, in 1763, crediting the bark of willow trees with a beneficial effect in alleviating distress due to fevers, aches, and pains. Later, the compound salicylic acid (named for the Latin word for willow, salix) was isolated from willow bark; it proved to be the active ingredient. By 1860, organic chemists were able to synthesize salicylic acid from basic starting materials, this furthered the therapeutic use of the substance, but there were problems. Salicylic acid proved to be irritating to the membranes of the throat, mouth, and stomach. These problems are directly associated with the high acidity of the compound, but a simple remedy was discovered, namely, replacement of the acidic phenolic hydrogen atom with an acetyl group.

When interpreting the structures of the above organic compounds, note the following characteristics of these molecules. Organic molecules are complex compounds of carbon. Carbon always bonds to four other groups or atoms. When outline structures (condensed) are given, such as the hexagon, each point of the hexagon represents a carbon atom. If a double bond, =, is present and the carbons are attached in a ring, only one hydrogen is attached to give the full compliment of four bonds. If a triple bond is present then only one other atom may be attached. Check the structures below to see that each carbon has four and only four bonds. Nitrogen, on the other hand, will bond covalently to only three atoms, and then oxygen bonds only to two.

A useful synthesis of acetylsalicylic acid was developed in 1893, patented in 1899, marketed under the trade name of “aspirin” by the Bayer Company in Germany. The name aspirin was invented by the chemist, Felix Hofmann, who originally synthesized acetylsalicylic acid for Bayer. More than 50 million 5-grain tablets of aspirin are consumed daily in the United States. In Part I of this experiment, you will prepare aspirin by reaction of salicylic acid with acetic anhydride, using concentrated sulfuric acid as a catalyst. By doing so you will learn some preparative procedures used in organic
synthesis in general. Aspirin still has its side effects, note that the carboxylic acid functional group remains intact. This may result in hemorrhaging of the stomach walls even with normal dosages. The acidic irritation can be reduced through the use of buffering agents, like antacids, in the form of magnesium hydroxide, magnesium carbonate, and aluminum glycinate when mixed with aspirin (Bufferin). While the ester can be formed from acetic acid and salicylic acid, a better preparative method uses acetic anhydrides in the reaction instead of acetic acid. An acid catalyst, like sulfuric acid or phosphoric acid, is used to speed up the process. In Part II of this experiment you will measure the percent acidity of aspirin by titration with a strong base.

Part I Synthesis of Aspirin
Caution! The preparation of aspirin involves the use of two hazardous materials – concentrated sulfuric acid and acetic anhydride. Proceed only if you have a fume hood to work in, and after you have listened carefully to the instructor’s safety directions.

1. Weigh 4.0 g (0.030 mole) of salicylic acid in a 125 mL Erlenmeyer flask. Using this quantity of salicylic acid to calculate the theoretical yield of aspirin. Record the weigh on the report sheet. 2. Carefully add 6 mL (0.051 mole) of acetic anhydride to the flask. (Care! Acetic anhydride is irritating to the skin and eyes.) 3. Using extreme caution, add 5 drops of concentrated sulfuric acid to the flask, swirl gently, and place the flask in a beaker of boiling water. Clamp the flask to a ring stand and heat for 20 minutes. Constantly stir with a glass rod; the entire solid must completely dissolve. 4. Remove the flask from the boiling water bath and allow to cool to room temperature. Crystallization should occur during cooling. If crystals begin to grow, let the flask sit undisturbed until crystals stop growing then add the 40 mL of ice water. If crystals do not grow, slowly pour the solution into a 250-mL beaker containing 40 mL of ice water, mix thoroughly, and place the beaker in ice water and let sit undisturbed until crystals have grown. The water destroys any unreacted acetic anhydride and will cause the insoluble aspirin to precipitate out of
solution. 5. Collect the crystals by vacuum filtration.

6. Wash the crystals with two 10-mL portions of cold water followed by one 10- mL portion of cold ethanol. Allow the crude product to dry, then weigh it on the rough balance. 7. Weigh a watch glass. Add the crystals and re-weigh. Calculate the weight of crude aspirin. Determine the percent yield. Test a small amount of this crude product for its melting point as described in Part II. Test the freshly made product for purity. Aspirin naturally decomposes into acetic acid over time so the purity test should be done the day the aspirin is prepared. Save some of your aspirin for testing.

The crude aspirin needs to be further purified. The crude products obtained from most preparations of organic compounds are contaminated with unreacted starting materials and substances from side-reactions. These can often be eliminated by a simple process known as RECRYSTAlLIZATION. The next phase of this experiment involves the recrystallization, and thus, purification, of your crude aspirin sample. 8. Dissolve about 2-4 g of your crude product in about 20 mL ethyl alcohol in a 125 mL Erlenmeyer flask, warming the alcohol in a water bath to effect dissolution. Caution: Do not use a flame to heat ethyl alcohol. As it is a flammable compound. If you obtained less than 6 g of crude product, use proportionately less alcohol. 9. If any solid material remains undissolved, filter the solution. 10. Add 50 mL of warm water (about 50oC) to the clear alcohol solution. If any crystals appear at this point, heat the contents of the flask until they dissolve. 11. Set the flask aside to cool, observing it carefully.

12. When crystals start to form, cool the flask by surrounding it with cold water. The crystallization process will then go to completion. 13. Collect the crystals by vacuum filtration.
14. Allow the crystals to dry.
15. Save your sample of the aspirin for a melting point determination and further analysis. Part II ANALYSIS OF ASPIRIN
A. Determination of the Melting Point
Most organic compounds have a sharp melting point, which can be measured
accurately to within 1oC or better, using the method below. Furthermore, the measurement is easily made with a small quantity of material (a few small crystals) using a simple apparatus. Melting point determinations are routinely made on solid organic compounds, and extensive compilations of melting points are available in reference books. One use of the melting point is to establish that a preparative or isolation procedure has led to an expected product. (The prepared substance should have the documented melting point for that substance.)

A very pure substance has a very sharp melting point. Further purification will not change the melting point. Less pure substances melt over a range of temperatures that is below the actual melting point of pure material. Thus the sharpness of a melting point is a useful criterion of purity. When a melting point is determined, it is therefore important that the melting range be recorded. The bottom of the melting range is the temperature at which the first signs of liquid can be seen. The top of the melting range is the temperature at which the last of the solid fuses, i.e. turns into liquid. The compound is generally regarded as pure enough for most purposes if the melting range is no greater than 2oC. A wide melting range signals the need for further purification.

1. Obtain a capillary tube from your instructor, and gently press the open end into the pile of aspirin crystals on the paper so that a few crystals of aspirin enter the capillary tube. 2. Tap the closed end of the capillary onto the bench top, so that the aspirin crystals work their way to the bottom. The aspirin crystals should be firmly packed, and fill the capillary tube to a depth of no more than 1-2 mm. Insert the capillary tube containing the sample into the melting point apparatus. Record the temperature where the melting point is first observed and when it becomes a liquid completely. This is your melting point range. Melting point of purified aspirin is 135-136 oC.

B. Determination of Purity
Phenols form a colored complex with the ferric ion. If phenol is present in
a sample, the resulting color means the product is impure. The purple color indicates the presence of a phenol group. The intensity of the color qualitatively tells how much impurity is present. Procedure

1. Label three test tubes; place a few crystals of salicylic acid into test tube #1, a small sample of your aspirin into test tube #2, and a small sample of crushed commercial aspirin into #3. Add 5 mL of deionized water to each test tube and swirl to dissolve the crystals. 2. Add 10 drops of 1% ferric chloride to each test tube.

3. Compare and record your observations.

Chemistry 51 ………………………Report Sheet – Aspirin ………………………..Experiment

1. Theoretical yield:
__________g (1 mol / 138 g salicylic acid) (180 g/ 1 mol) = _________________g of aspirin Mass of salicylic acid Theoretical yield of aspirin

2. Experimental yield:
Weight of aspirin & watchglass________________g
Weight of watch glass________________g
Mass of crude product obtained after suction filtration________________g Percent Yield of crude product (expt / theo * 100)________________%

Mass of re-crystallized product (optional)________________g Percent Yield of re-crystallized product (expt / theo * 100)________________% Part II
Melting Point of crude product (1st trial)______ (2nd trial)_______

Melting Point of re-crystallized product (1st trial)______ (2nd trial)_______

Ferric Chloride Test (Purity test)

Salicylic acid

Your aspirin

Commercial aspirin

1. What is the purpose of the 18M sulfuric acid in the preparation of aspirin?

2. Explain why the percent of your aspirin was different from the results obtained from the commercial aspirin.

3. Old aspirin exposed to moisture often smells like acetic acid (vinegar). When aspirin is heated in boiling water, it decomposes and gives off a vinegar smell. The resulting solution gives a positive FeCl3 test. Why is this test positive? Write the chemical equation for the reaction of aspirin and water at high temperatures.

4. You have spent about 15 weeks learning about chemistry and its respective laboratory techniques. The typical street-drug producer usually doesn’t have any chemistry background but instead simply mimics the synthesis learned from another person or, in some cases, a book. If you take this into consideration along with your success at producing aspirin, how well do you trust the producers of street drugs to make drugs safe for consumer use?

Cite this Aspirin lab Essay

Aspirin lab Essay. (2016, May 11). Retrieved from

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