Introduction Before the synthetic aspirin of today, salicylic acid, which is the important ingredient found in aspirin, was extracted naturally from methyl salicylate found in Wintergreen oil, which could be found in certain plants. The purpose of this lab experiment is to prepare salicylic acid from the natural starting material, methyl salicylate, and compare it with the salicylic acid produced from the artificial compound benzene (obtained through petroleum refining). Based on the two reactions of aspirin synthesis shown below, methyl salicylate and benzene share a common intermediate which is salicylic acid:
In order to prepare salicylic acid from wintergreen oil, an organic synthesis must be carried out. This is done first by separating salicylic acid from methyl salicylate, followed by its purification. The separation can be done by performing a reflux of a solution of methyl salicylate and sodium hydroxide. After the reflux is complete, sulphuric acid will be added to the refluxed solution, which will cause the salicylic acid to precipitate. This will then be filtered through suction filtration.
Once the product has dried, purification must be done by recrystallization; dissolving the product in a minimum amount of hot solvent (water), and then filtering once again through suction filtration. Finally, to complete the main purpose of the experiment, which is to compare the prepared product with a standard (salicylic acid from benzene), an analysis must be done by measuring the melting point of the crystals and of the standard.
This is fulfilled by using capillaries opened at one end containing a very small amount of both samples (dried product and a 1:1 mixture of the dried product with standard), and inserting them into the melting-point apparatus, where their melting point can be read on the thermometer.
Results: Physical Properties: | Molecular weight (g/mol)| Melting point (oC)| Boiling point (oC)| Density (g/mL)| Methyl salicylate| 152. 1| -8| 223| 1. 174| Salicylic acid| 138. 1| 159| -| -| Raw data: Mass of watch-glass (g)| Mass of watch-glass + initial product (g)| Mass of watch-glass + crystals (g)| Melting point range of crystals (oC)| Melting point range of 1:1 mixture (oC)| 29. 677| 30. 808| 30. 876| 154-157| 158-159|.
Observations: Precipitation of salicylic acid occurring from addition of sulphuric acid crystals appearing on the surface while cooling in ice-bath Data: Mass of product, actual yield (g)| Percent yield (%)| Melting point range (oC)| 1. 199| 86. 82| 154-157| Calculations: Theoretical yield *Methyl salicylate =MS, Salicylic acid= SA 1. 521g MS x (1mol MS/152. 1g MS) x (1mol SA/1mol MS) = 0. 01 mol SA Mass of SA: 0. 01mol SA x (138. 1g SA/1mol SA) = 1. 381g Salicylic acid Percent yield Percent yield = (actual yield/ theoretical yield) x 100% = (1. 199g/ 1.
381g) x 100% = 86. 82 % Percentage error Percentage error = [(Actual yield – theoretical yield)/ theoretical yield] x 100% = [(1. 199g – 1. 381g)/ 1. 381g] x 100% = 13. 18% Mechanism: Discussion and conclusion Methyl salicylate (which comes from wintergreen oil) and benzene are two compounds from which salicylic acid can be obtained from. The difference between the two is that one can be found naturally in plants (methyl salicylate) and the other is artificially produced (benzene). The question is do they yield the same desired product (salicylic acid)?
This lab experiment consisted of preparing salicylic acid from a methyl salicylate sample and, and comparing it with salicylic acid made from benzene by measuring their melting points at the same time. In order to do so, the first step was to mix about 1. 296mL of methyl salicylate with 15mL of 6M sodium hydroxide in a round bottom flask. This solution will undergo reflux for about half an hour. The reflux is done in order to heat the mixture without losing any of the product and to increase the reaction rate.
By adding the sodium hydroxide, which is in excess, this will form an intermediate anion; sodium salicylate. This is an acid-base reaction; the sodium salicylate that is formed is water-soluble and is a weak base. The goal of this lab is to produce salicylic acid, so in order to do so, a strong acid must be added to neutralize the insoluble salicylic acid; sulphuric acid. Therefore, once the reflux is done and the flask has cooled down, 16mL of 3M sulphuric acid was added. By adding the strong acid, this has caused the salicylic acid to precipitate.
The pH was then tested to make sure it was around 2. By testing the pH, it is possible to assure that all of the salicylic acid has been neutralized since the pKa of salicylic acid as about 3; so if the pH is above 2, more acid would have to be added to bring it down. This mixture was then cooled down in an ice-bath for about 10 minutes. This was then followed by a suction filtration in order to obtain the solid and dry it. A small amount of cold water was used to wash the solid to get rid of impurities. This solid was then weighed (1.131g). Next, in order to obtain a much purer sample, a recrystallization had to be made.
This is done by dissolving the salicylic acid in a minimum amount of boiling distilled water. In this step it is very important not to add too much solvent, other ways the maximum yield of the product cannot be obtained. Just a small amount of hot solvent (enough to completely dissolve the product) must be added to get rid of impurities present. Once the solid is dissolved, this was then filtered by vacuum filtration to obtain the crystals.
The crystals were then washed (to dissolve still existing impurities and get rid of them) and then dried, and finally weighed (1. 199g). Ice-cold water was used in the washing process in order to prevent some of the crystals to completely dissolve in it and be lost. The last part of the experiment consisted of analysing the crystals by measuring its melting point range and comparing it with a standard. This was done by inserting about 1cm of the crystals into a capillary (open at one end), and the same amount of a 1:1 mixture of the crystals with the standard into a separate capillary.
Both capillaries were then inserted into the melting point apparatus, where their melting point ranges were measured at the same time; crystals have a melting point range of 154-157oC and the 1:1 mixture, 158-159oC. The melting point ranges that were measured show that salicylic acid has indeed been synthesized, but that the fact that the range is slightly lower than that of the standard confirms a small presence of impurities. The percent yield was found to be 86. 82%, and the percentage error, 13. 18% .
The causes of error are the reason for this yield; a loss of product could have occurred during the transferring process of the solution from the boiling flask to the Buchner funnel in the first filtering step. The same thing might have happened during the separation of the precipitate from the liquid. Another source that can explain the loss of product is the addition of too much hot solvent during the recrystallization process; some of the salicylic acid would have completely dissolved so that it cannot be retrieved after filtration.
One way to improve and maximize the yield is to always wash the flask or beaker that contained the compound after transferring it into the Buchner funnel with cold distilled water to ensure that no product is left. Finally, based on the analysis of the crystals, the objective of preparing salicylic acid has been met. Salicylic acid has been synthesized from methyl salicylate using 3 major techniques: reflux, suction filtration and recrystallization. To confirm that salicylic acid was synthesized, the melting point range had to be measured, and was found to be very close to the 1:1 mixture’s melting point.
References Lehman, John W. Operational Organic Chemistry, 4th ed. ; Prentice Hall, 2008. p. 61-68 Lehman, John W. Operational Organic Chemistry, 4th ed. ; Prentice Hall, 2008. p. 843-846 Fleming, Steven A. ; Jones, Maitland Jr. Organic Chemistry, 4th ed. ; New York, 2010. Smith, Michael B. ; March, Jerry. Organic Chemistry, 6th ed. ; New Jersey, 2007. Laurence M. Harwood, Christopher J. Moody, Jonathan M. Percy. Experimental organic chemistry: standard and microscaling. ; 1998.