This report gives a detailed account of the experimental preparation of aspirin in a laboratory environment as well as some basic industrial background on the product. It also contains information about safety precautions put in place to ensure the safety of the team who carried out the experiment. All results obtained have been included as well as a detailed analysis of what they represent alongside any improvements to the method used. Introduction.
The aim of this experiment was to successfully prepare a sample of 2-acetoxybenzoic acid [1], commonly known as aspirin, through the reaction of salicylic acid with acetic anhydride under acidic conditions provided by sulphuric acid. Once the reaction was completed the samples’ yield and melting point were determined for comparison to industrial perimeters. The equation below outlines the overall reaction between salicylic acid and acetic anhydride. Note that acetyl salicylate is a synonym of aspirin and acetic acid is a by-product.
Aspirin is one of the safest pain relievers on the market, its great success as a non-prescription drug also comes down to its affordable price as well as its other medicinal features such as having an anti-inflammatory affect. 80 million aspirin tablets [3] are consumed annually in the US alone therefore it must be manufactured at a grand scale. The solubility of aspirin in water at 20°C is 3mg/ml [4]; however this value can change with a change in temperature.
[5] At higher temperatures aspirin molecules have a higher kinetic energy and therefore collide with each other as well as water molecules more often, resulting in an increase in solubility. On the other hand, at lower temperatures aspirin molecules have lower kinetic energy and this allows them to bind together, also known as ‘falling out’ of solution or crystallisation.
The main variations between the method used in the experiment and the common industrial methods of synthesising aspirin are; the equipment used (e. g. large 1500 gallon, glass-lined reactor tanks, agitators and water-cooled reflux condensers), toluene being used as a solvent which does not take part in the reaction, time given for the reaction to take place (normally around 24 hours), cool down period of the mixture (around 3-4 days), and the temperatures at which the reaction is carried out (88-94 degrees Celsius) [6].
Method The method for carrying out the experiment successfully involved a few simple steps some of which had to be continuously over a long period of time. Firstly 50g of salicylic acid is mixed with an 80ml solution of acetic anhydride contained within a round-bottomed Quick-fit flask, which had been previously set up with a water-cooled condenser and a fitted thermometer.
After this 0. 5ml of concentrated sulphuric acid was measured out added to the solution followed by a stopper to prevent the escape of potentially dangerous fumes. For the next 30 minutes, the flask was gently heated at 70°C and every 2 minutes agitated to ensure that the sulphuric acid was evenly dispersed allowing the salicylic acid to dissolve into the solution and therefore react. Whilst agitating the solution, an occasional large increase in temperature occurred which had to be countered by lowering the heat source away from the flask in order to prevent the temperature from exceeding 70°C.
The time and temperature were monitored using a stopwatch and the thermometer respectively. Once the 30 minutes had passed, the flask was allowed enough time to cool to 50°C. The contents of the flask were then poured out into a beaker containing 750ml of cold distilled water which acted as an ice bath. After allowing time for the acetic acid to dissolve in the water, the contents were filtered out into an Erlenmeyer flask using a funnel lined with filter paper, where the insoluble aspirin collected on top of the filter paper.
Using a spatula, the aspirin was scrapped into a beaker which had been weighed previously. The aspirin was then given a week to dry before being weighed using the same balance used for weighing the empty beaker. A small sample of the aspirin was then collected in a glass tubule and placed into a melting point apparatus. The sample was carefully monitored until it completely melted, at which point the temperature was noted as that was the melting point of the aspirin. All results were collated in a table of results upon acquiring them. This concluded the practical section of the experiment. Hazards and Risk Analysis.
The chemicals used during the experiment all had potential safety risks associated them. In order to counteract these risks personal protection equipment was a non-negotiable requirement. This included splash goggles, latex gloves and a lab coat. Acetic anhydride is also known to be flammable and therefore it was kept away from any sources of ignition and this also meant that the heating apparatus for heating the flask was fully electric, where all electrics were grounded. The entire experiment was carried out in a fume cabinet in order to prevent the inhalation of any harmful fumes produced during the reaction.
In case of a spillage the spilled liquid would have to be soaked up by an inert solid and placed in the suitable waste disposal bin. In case of skin or eye contact the area must be thoroughly washed for 15 minutes whilst all contaminated clothing is removed. For a more detailed insight into all safety precautions please refer to the MSDS forms of all chemicals used in and formed from the experiment, enclosed with the report. Results and Calculations Item Mass / g Mass of Empty Beaker = 175. 96 Mass of Beaker and Aspirin = 223. 20 Resultant Mass of Aspirin = 47. 24 Melting Point of Aspirin = 122°C Calculating Percentage Yield of Aspirin.
The resultant mass of aspirin = mass of beaker and aspirin – mass of empty beaker = 223. 20g – 175. 96g = 47. 24g In order to calculate the percentage yield of aspirin the limiting reagent of the reaction was identified first. Moles of Salicylic Acid Mass of Salicylic Acid = 50g Mr of Salicylic Acid = 138. 1g/mol Therefore Moles of Salicylic Acid = Mass/Mr = 50g/138. 1g/mol = 0. 362 moles Moles of Acetic Anhydride Mass of Acetic Anhydride = volume x density Density of Acetic Anhydride = 1. 08g/cm3 Therefore Mass of Acetic Anhydride = 80ml x 1. 08g/cm3 = 86. 4g Mr of Acetic Anhydride = 102. 1 g/mol Moles of Acetic Anhydride = 86. 4g/102.
1g/mol = 0. 846 moles Salicylic acid is the limiting reagent as it has fewer moles than acetic anhydride therefore the number of moles of aspirin produced is equal to the number of moles of salicylic acid, 0. 362 moles. Theoretical Yield of Aspirin Theoretical yield = moles x Mr of product Moles of Aspirin = 0. 362 moles Mr of Aspirin = 180. 2 g/mol Therefore theoretical yield of aspirin = 0. 362mol x 180. 2g/mol = 65. 23g The actual yield of aspirin was 47. 24g. Percentage Yield of Aspirin Theoretical yield of aspirin = 65. 23g Actual yield of aspirin = 47. 24g Percentage yield = (actual yield x 100%) / theoretical yield = (47.24g x 100%) / 65. 23g = 72. 42% .
Discussion and Recommendations From the results obtained it can be concluded that the sample of aspirin produced through this experiment contained some impurities which resulted in a melting point of 122°C. This was a lower value than that of the generally recognised melting point of aspirin which is thought to be 135°C [7]. The impurities must partially consist of the reagents used in the reaction as the percentage yield indicates that only 72. 42% of the theoretical amount of aspirin was produced. This means that the reaction was incomplete and some reagents remained in the solution.
However the most likely reagent to cause the impurity is salicylic acid as it is insoluble in water and would therefore be filtered out with the aspirin. Another possible impurity could be the acetic acid, as it may not have had enough time to fully dissolve into the ice bath which was used to crystallise the aspirin molecules. A significant quantity of aspirin was also left in the funnel after filtering which meant that the actual yield of aspirin which was determined through weighing was actually slightly lower than it should have been, this would therefore have resulted in a lower percentage yield.
There are a few ways of improving the experiment carried out in order to increase the reliability as well as the validity of the results. One way of improving the experiment is by using an actual agitator to agitate the contents of the flask. This would ensure a more even distribution of sulphuric acid as well as being a safer method for agitation. Another improvement would be to use a more accurately set heat source when heating the flask to further reduce the amount of times the temperature exceeded 70°C.
When scrapping out the aspirin from the funnel instead of using a spatula, the use of a fine brush would allow for more aspirin to be obtained. As a final remark it can be said that the experiment was an overall success as a considerable amount of aspirin was produced. Nomenclature List with Units g – Unit of mass which stands for grams. °C – Unit of temperature which stands for degrees Celsius. Mr – Abbreviation for relative molecular weight. mol – Unit for the amount of a substance which stands for moles. g/mol – Unit of relative molecular weight which stands for grams per mole. ml – Unit of volume which stands for millilitres.g/cm3 – Unit of density which stands for grams per centimetre cubed. mg/ml – Unit of solubility which stands for milligram per millilitre.
References and Acknowledgements [1] – International Union of Pure and Applied Chemistry (IUPAC) [2] – http://www. britannica. com/EBchecked/topic/519177/salicylic-acid [3] – http://www. madehow. com/Volume-1/Aspirin. html [4] – Pharmacokinetic Solubility And Dissolution Profile of Non-Steroidal Anti-Inflammatory Drugs [5] – http://chemistry. about. com/od/demonstrationsexperiments/ss/aspirin_5. htm [6] – http://njcmr. njit. edu/distils/lab/aspirins/nap4. html [7] – Handbook of Chemistry and Physics.