Synthesis of Aspirin Lab

Synthesis of Aspirin

Worked with Grace Choi

  1. Introduction

In status quo, aspirin is a commonly bought, widely used over-the-counter drug. It has lots of functions: it can reduce fever, pain, swelling, soreness and redness. Historically, its first account was by a Greek physician called Hippocrates when he recorded the effects of the powder created from the willow tree at curing childbirth and fever. Then, its active ingredient, salicylic acid was first discovered by an English pastor, Reverend Edward Stone, when he administered powdered willow bark on patient and observed its effect on them. Friedrich Kolbe then successfully described both structure of salicylic acid and its method to synthesize it. Salicylic acid was used to cure headache and pain since then. However, for its high acidity, salicylic acid could not be prescribed for long. Felix Hoffman then synthesized acetylsalicylic acid to activate the effects when decreasing the irritation in body, and named it aspirin. (Szczeklik 1) Proved by history, aspirin has lots of pharmacological benefits for curing illness. In this experiment, aspirin was synthesized with salicylic acid and acetic anhydrate.

Aspirin is synthesized through a mechanism, now called aspirin synthesis mechanism. For the sake of time and efficiency, sulfuric acid was added as a catalyst to speed up the reaction. In the process of aspirin synthesis mechanism, the protons of the acid reacts with the acetic anhydrate to bond with both oxygen with double bond and in between the acetate groups. This will create positive charge on both the protonated oxygen. Thus, the oxygen will hog the electrons from the C=O double bond to move its positive charge to carbon. Salicylic acid is consisted of a phenol attached to a carboxylic acid group. The oxygen in phenol, being more electro negative than carbon will be slightly negative than its surrounding. This will attract the positive carbon creating a bond with it. To release the positive charge on oxygen, the deprotonated acid will take away H+. Then, the protonated oxygen in between the acetate group, having a positive charge, will hog electrons from the C-O bond that it has with the carbon nearer to the phenol group. This will break the bond between oxygen and carbon creating acetic acid. Due to this, now carbon has a positive charge which will attract the lone pairs of oxygen creating a double bond, yet again, creating positive charge on oxygen. However, the second deprotonated acid will take away H+ from the oxygen to minimize formal charge. The final product is the phenol group attached to a carboxylic group and an acetate group, aceticsalicylic acid, aspirin, and acetic acid. To visualize the mechanism, a figure below can be used.

Aspirin synthesis mechanism

Figure 1

When the synthesis was over and the crystals formed, water was added to destroy all unreacted acetic anhydride.

Both salicylic acid and aspirin are not very polar. Thus, they are slightly soluble in water when being highly soluble in ether or other organic compounds. (Rainsford 77, 78) Thus, acetic anhydrous, being an ether and an organic solvent that can dissolve salicylic acid and aspirin easily, was used as the solvent in which the reaction took place in. However, due to their large molecular mass, salicylic acid and aspirin has limited solubility meaning it is still not highly soluble as common salts such as NaCl. Due to this property, the solubility of both salicylic acid and aspirin depends largely on the temperature of the solvent. Thus, cold water and ethanol was used to rinse the solution and the solvents in which salicylic acid and aspirin needed to be dissolved was placed in the warm bath to increase temperature.

Recrystallization is a process which increases purity of crystallized solid product. When aspirin crystals are collected, there are possibilities that impurities that are still soluble in polar solvent remain in solid form. Thus, the solid product was dissolved in alcohol, a polar solvent, with heat. Then, the crystals were collected again.

To check for impurities that might remain in the product, titration and back titration method was used. Titration is a laboratory procedure that is used to determine the concentration of an unknown solution using a solution of known concentration. It is done by slowly adding the acid or base solution of known concentration to the solution of unknown concentration to look for the point where all substances are neutralized. Similarly, back titration is a process that is used to ensure the process of titration. It is carried out by slowly adding the solution of opposite acidity of known concentration to the solution that is once titrated and is added with excess of the first titrant in order to look for the point where the solution is neutralized completely. Both salicylic acid and aspirin, containing a carboxylic acid group, is acidic in nature and they both, along with other acid impurities, may exist in the final product. However, aspirin out of all compounds not only can deprotonate using NaOH but also can use NaOH to turn itself back into salicylic acid through a slow process called hydrolysis as shown in figure 2. As shown in figure 2, aspirin reacts with NaOH in 1:1 ratio to produce salicylic acid. Thus, after the first titration in which all acidic compounds are deprotonated, more NaOH was added to react with aspirin to form hydrolyzed aspirin, salicylic acid. When all aspirin has reacted with NaOH to hydrolyze, HCl was added to react with excess NaOH. Therefore, at the point of neutralization, (the moles of excess NaOH added after the first titration) – (the moles of HCl added for back titration) equals to the moles of NaOH reacted with aspirin to produce salicylic acid. This value also equals to the amount of aspirin that existed in the product in the first place since aspirin did not hydrolyze during the first titration, just deprotonated. Therefore, to obtain the purity by percent, the following equation can be used.

Titration of Aspirin

Figure 2

  1. Apparatus
    1. Salicylic Acid
    2. 125mL Erlenmeyer Flask
  • Acetic Anhydride
  1. 16M Sulfuric acid
  2. 500mL beaker
  3. Heating plate
  • Stand
  • Clamp
  1. Pipet
  2. Cold water and ethanol prepared in fridge
  3. 250mL beaker
  • Vacuum Filtration set
    • 500mL Erlenmeyer flask with side arm
    • Erlenmeyer flask stopper
    • Filter paper
    • Pump
    • Rubber cable
    • Glass cylinder funnel
    • Glass support base with sand core permeable top
    • Metal clamp
  • Weighing boat
  • Phenolphthalein Indicator
  1. Distilled Water
  • 0.1M NaOH
  • 0.1M HCl
  • Buret
  • Magnetic Stirrer
  1. Procedure
    1. Synthesis
      1. Weigh 40 of salicylic acid in a 125mL Erlenmeyer Flask. Record the exact mass.
      2. Add 6mL of acetic anhydride to the flask
  • Add 5 drops of sulfuric acid to the flask, swirl gently, and place the flask in a beaker of boiling water. Clamp the flask to a stand and heat for 20 minutes. Stir with glass rid, or if not present, glass pipet. Solid must all dissolve
  1. Remove the flask from the water bath and cool in room temperature. It should crystalize.
    • If crystals grow, let them grow until it stops. Pour 40mL of cold water
    • If not, pour the solution into a 250mL beaker of 40mL of cold water, mix thoroughly.
      • Water will destroy unreacted acetic anhydride and insoluble aspirin will precipitate in solution
  1. Collect the crystal by vacuum filtration.
  2. Rinse the crystal with 10mL of cold water and 10mL of Ethanol, let it dry and weigh it.
  1. Recrystallization
    1. Dissolve your crude product in about 20mL ethyl alcohol in a 125mL Erlenmeyer flask. Warm the alcohol in a warm bath to speed up dissolution.
    2. If any solid remains, filter it.
  • Add 50mL of warm water to clear alcohol solution.
  1. Keep heating to dissolve it.
  2. Set the flask aside to cool.
  3. If crystals form, cool the flask by surrounding it with cold water.
  • Collect the crystals by vacuum filtration.
  • Allow the crystals to dry.
  1. Save the sample.
  1. Analysis by Titration
    1. Measure out 0.5g of the synthesized Aspirin.
    2. Dissolve it in 15mL of distilled water. Add 5 drops of Phenolphthalein Indicator.
  • Use a stock solution of 0.1M NaOH solution to titrate the aspirin solution
  1. When the indicator turned color, add twice as much of the NaOH used for the titration to the solution. Add as much of the NaOH as the Aspirin solution first acquired (15mL).
  2. Heat the solution in warm bath for 10 minutes.
  3. Back titrate the solution using 0.1M of HCl.
  • When the solution turned color, test for the accuracy by adding drops of NaOH and HCl to titrate and back and forth. If the titration was accurate, the solution will change color with a drop of NaOH or HCl.
  1. Result
    1. Synthesis
      1. Mass of salicyclic acid + weighing boat
        • 807
      2. Mass of weighing boat
        • 777
  • Mass of Salicyclic acid
    • 030g
  1. Mass of Weighing boat
    • 781g
  2. Mass of Aspirin and weighing boat
    • 629g
  3. Mass of Aspirin
    • 048g
  4. Recrystallization of Aspirin
    1. Mass of filter paper
      • 785g
    2. Mass of filter paper + Recrystallized Aspirin
      • 069g
  • Mass of recrystallized aspirin
    • 284g
  1. Analysis by titration
    1. Mass of Aspirin used for titration
      • 503g
    2. Volume of NaOH used for first titration
      • 24mL
  • Additional NaOH included
    • 24 + 15 (mL)
    • 39mL
  1. Total volume of NaOH used
    • 24 * 24 + 15 (mL)
    • 63mL
  2. Volume of HCl used for back titration
    • 5mL
  3. Analysis
    1. Purity test
      1. Mass of Aspirin first created
        • 048g
      2. Mass of Aspirin after recrystallization
        • 284g
  • Mass of impurity discarded by recrystallization
    • 764g
  1. Percent purity by titrationpurity calculation
  2. Test for purity
    • 39mL*0.1M*1L/1000mL – 15.5mL*0.1M*1L/1000mL = 0.00235moles = total moles of aspirin
    • 00235moles*180.16g/mol = 0.423g = total mass of aspirin that is present in product
    • (Total mass of product) – (Total mass of aspirin) = (Total mass of impurities)
      • 503g – 0.423g = 0.08g
    • (Total mole of NaOH used to titrate all acidic compound of solution) – (Total mole of NaOH used to hydrolyze aspirin) = (Total mole of NaOH used to titrate impurity)
      • 24mL*0.1M*1L/1000mL – 23.5mL*0.1M*1L/1000mL = 0.5mL*0.1M*1L/1000mL = 0.00005mole of NaOH used to titrate impurity = Total mole of impurity assuming 1:1 reaction ratio
    • 08g/0.00005mole = 1600g/mol = average molar mass of impurity assuming full reaction and 1:1 reaction ratio
  3. Discussion

The percent purity of aspirin turned out to be 97.91%. It was a very possible value considering the recrystallization of the crude aspirin product obtained after the first synthesis. However, to discuss about the impurities, some compounds created and destroyed in the synthesis process can be taken into consideration.

One possibility is the salicylic acid. Even though the mechanism of synthesis of aspirin is defined, it cannot be said that salicylic acid has reacted fully to form aspirin because of the complexity of the mechanism. Thus, some salicylic acid would have been left in the solid crystals to react with NaOH when titrating.

Another possibility is the acetic acid. In the process of synthesis of aspirin, acetic acid is also produced. Furthermore, when water is added to unreacted acetic anhydride, acetic acid is formed. Even though, acetic acid is highly miscible with water, there is a slight possibility that it could have stayed with the final product to be titrated.

Furthermore, in the process of cooling, the solution containing the precipitate was kept in a fridge, where a mixture of lead nitrate and potassium iodide also was kept. Staying in the fridge for a significant amount of time might have resulted in some lead particles being transferred to the aspirin solution. This would explain the ridiculously high average molecular mass of the impurity. Since lead ions would not react with acid, the total mass of compounds that took part in the titration reaction would significantly decrease resulting in a more reasonable average molecular mass of the impurity.

To improve the experiment, the precipitates should have been washed carefully to get rid of the water soluble impurities, in this case, acetic acid. Furthermore, the fridge in which the experiment materials are left should be clean from other compounds. This improvement will increase accuracy by not adding extra mass to the precipitate that will not react when it is assumed to. Additionally, since this experiment dealt with fine powder of compounds, the scientists have to be very careful at handling and transferring of the powders from one place to the other.

  1. Reference
    1. Resources
      1. Kemper, J. Aspirin Titration
      2. Los Angeles City College. Synthesis and Analysis of Aspirin. 2005
  • Rainsford, K. D. Aspirin and Related Drugs; CNC Press: New York, 2004.
  1. Szczeklik, A. The History of Aspirin: The Discoveries that Changed Contemporary Medicine. Paths of Discovery 2006, 18, 175-184
  1. Figures
    1. Duldren. A mechanism describing the formation of acetylsalicylic acid (Aspirin) from acetic anhydride and salicylic acid. Chemdraw 2012
    2. Kemper, J. Aspirin Titration


For the Actual Document, click Synthesis of Aspirin.


Leave a Reply

Fill in your details below or click an icon to log in: Logo

You are commenting using your account. Log Out /  Change )

Google+ photo

You are commenting using your Google+ account. Log Out /  Change )

Twitter picture

You are commenting using your Twitter account. Log Out /  Change )

Facebook photo

You are commenting using your Facebook account. Log Out /  Change )


Connecting to %s