Introduction

Almost all industrially-synthesised chemicals require purification, a separate process that takes extra time, money, space and expertise. The initial aim was to design a more efficient installation, condensing the purification process into the synthesis to obtain a pure product without any further processing. This study was conducted on the industrial production of acetylsalicylic acid as neither the reactants nor products are toxic or carcinogenic; the reaction temperature and pressure can be easily reached in lab and the reaction is quick so that many experiments can be made in a short time. Furthermore the reaction used in-lab is the same as the the reaction used industrially 1 .

In the case of acetylsalicylic acid, the purification phase is recrystallization, which involves dissolving the impure product in a warm solvent and then cooling the solution slowly. In this way, the product of interest precipitates, leaving all the impurities dissolved. Following simple filtration, the product is purified. However, in order for the two steps to be performed together, the solvent needs to have some “extra” features:

  • It must withstand the reaction conditions (it mustn’t boil, polymerize or decompose)
  • It must not react with any of the reactants or the products
  • It must be cheap and possibly ecologically friendly

A preliminary study was carried out to choose the right solvent. A further study was then made to determine the right quantity of solvent as this affected the time taken by the whole process.

Literature Review

A number of extremely efficient processes already exist for the production of acetylsalicylic acid. These involve the use of catalysts 2 which speed up the reaction without affecting its outcome. However, the catalyst itself is lost in the batch and cannot be recovered; no studies seem to tackle the problem of purification itself.

Method

Finding a suitable solvent

The first series of experiments were carried out to identify a suitable solvent. Alcohols and water were excluded as they would have hindered the reaction. Alkanes and alkenes were excluded due to their low boiling points and difficulty of use; short-chained aldehydes, ketones, and carboxylic acids were left out too. Experiments carried out with formic acid showed that the reaction seemingly did not take place, perhaps due to the dehydration of the acid by the acetic anhydride.3 This reaction, competing with the esterification, might have vitiated the outcome. Another effect that might have played a role is the Le Chatelier principle – structural similarities or acid strength may have meant that the reaction was already in equilibrium. Experiments with acetone, however, yielded interesting results – simply evaporating the solvent resulted in extremely pure product.

The procedure used in all the experiments (similar to those described in the literature)4 , is the following:

5g of salicylic acid was added to 10 ml of acetic anhydride plus the chosen quantity of solvent in a 75 ml flask5. The mixture was then gently stirred into flask which was placed into a water bath at a temperature between 55°C and 60°C. Different experiments were carried out to explore the influence of open and closed systems. After fifteen minutes, the flask was placed to cool at room temperature, or under cold water. The product was then filtered by vacuum and washed with cold distilled water. After being dried, a few milligrams of product was placed in a capillary test tube, which was put in a melting point instrument, and a measurement was taken.

A known quantity (≈0.02 g) was weighed, put in a 100 ml volumetric flask and dissolved with 10 ml of denatured alcohol. Distilled water was added up to third of the flask. After mixing it until the solid was dissolved, 3 ml of 1% FeCl3 solution was added. The flask was filled with water and the whole solution was mixed again. The FeCl3 reacted with salicylic acid, the only pollutant left from the filtration, giving a strong purple colouration, which can be measured using a colorimeter. 6

Finding a suitable amount of solvent

Having completed the first series of experiments (which indicated acetone to be the most suitable solvent), further experiments were conducted to ascertain what the best proportion was. With 15 ml of solvent 7, it took a short time for the salicylic acid to dissolve though it took a night for the product to precipitate. The same thing had happened with 10 ml; it took a few hours for the product to precipitate when 5 ml was added. When 3 ml was added to the mixture, it took almost five minutes for the salicylic acid to dissolve, but the precipitation happened immediately as the mixture reached room temperature.

It was noted that the quantity of crystals precipitated depended on whether the system was closed or opened. When 1 ml of solvent was added to the solution, the salicylic acid did not dissolve completely within 15 minutes. The shape of the crystals was heavily influenced by the precipitation time which in turn depended on the quantity of solvent used. Finally, a control experiment was made adding neither any solvent nor any acid catalyst. The salicylic acid did not completely dissolve into acetic anhydride, which meant that the precipitate obtained at the end of the reaction was a mixture of both reactant and product.

Results

The precipitation time was measured from the end of the reaction though the measurement was of a subjective and purely observatory scope. The melting point of pure acetylsalicylic acid is 136°C8; in the last two experiments the reaction didn’t complete.

Quantity of Solvent (ml) System (open/closed) Precipitation time Nature of Crystals Melting point (°C)
15 Closed 24h Vitreous, needle like, 1-2 cm long 135.7
10 Closed 24h Vitreous, needle like, 1-2 cm long 135.6
5 Closed 18-24h Vitreous, needle like, 0.5-1 cm long 135.6
3 Closed 10-15 min 3mm cubes, white 135.5
3 Open 5 min White shiny dust 135.5
1 Open // // //
0 Open // // //

Photos

Results photo

Fig 1: 3ml of solvent with open system.

Results photo

Fig 2: 3ml of solvent with closed system. The shape of the crystals may be due to the slower cooling

Results photo

Fig 3: Results for 5 and 10 ml of solvent. The crystals are blue due to an indicator (Congo Red) used to highlight them

Analysis

The colorimetric test was only conducted on the 3 ml open system experiment as the product was most interesting – the crystals were extremely small and tended to form a paste-like suspension. This is advantageous as it is not only easy to move between containers, but is also very easy to filter. The results are shown below:

The standard was prepared with 0.0259 g of salicylic acid diluted using water in a ratio of 1:10 (it would otherwise be opaque).

Type Weight (g) Absorbance
Standard 0.0026 0.229
Sample 0.2140 0.017

The equation of the line of best fit is Y= 88.077X. Solving it for the Y (ABS) value 0.017 yields 2×10-4 g of impurities after washing and filtration the product indicating a purity of:

Purity= ==99.9065%

Conclusions

S1: Acetic anhydride reservoir S2: Acetone reservoir S3: Salicylic acid reservoir R1: Reactor D1: Distillation column B1: Continuous oven, for drying the salicylic acid. P1: Feed pump relative to S1 P2: Feed pump relative to S2 P3: Feed pump relative to S3 P4: Feed pump for acetone recovery P5: Feed pump for salicylc acid recovery P6: Feed pump for washing water (recovered from distillation) P7: Feed pump for acetic acid SC1: Heat exchanger for the mixture SC2: Heat exchanger for acetone SC3: Heat exchanger for the washing water SC4: Heat exchanger for acetic acid F1: Rotary drum filter for acetylsalicylic acid recovery F2: Rotary drum filter for salicylic acid recovery

  • S1: Acetic anhydride reservoir
  • S2: Acetone reservoir
  • S3: Salicylic acid reservoir
  • R1: Reactor
  • D1: Distillation column
  • B1: Continuous oven, for drying the salicylic acid.
  • P1: Feed pump relative to S1
  • P2: Feed pump relative to S2
  • P3: Feed pump relative to S3
  • P4: Feed pump for acetone recovery
  • P5: Feed pump for salicylic acid recovery
  • P6: Feed pump for washing water (recovered from distillation)
  • P7: Feed pump for acetic acid
  • SC1: Heat exchanger for the mixture
  • SC2: Heat exchanger for acetone
  • SC3: Heat exchanger for the washing water
  • SC4: Heat exchanger for acetic acid
  • F1: Rotary drum filter for acetylsalicylic acid recovery
  • F2: Rotary drum filter for salicylic acid recovery

The installation shown here is still a just a sketch. Some wasted energy can be reused in the installation itself (the steam obtained during the distillation can be used to heat the reactor; the hot products can be used to pre-heat the reactants, cooling themselves, etc). No calculations about the liquid fluxes or the energy were made; however, it still shows how the results can be easily applied to an industrial context.

The results show that is possible to synthesise pure products from the synthesis environment with a high yield, between 60% and 70% . The experiments have demonstrated the crucial role of the solvent, indicating that it is necessary to dissolve the products to take them to a reactive state (though it keeps the acetylsalicylic acid dissolved, reducing the yield). This led to the idea of opening the system right after the salicylic acid is dissolved. In this way, much of the acetone to evaporates which is necessary in order to keep the impurities dissolved in solution.

In addition, when the product is washed with water the acetic anhydride in the remaining solution undergoes hydrolysis, which produces acetic acid. The remaining solid, which is a mixture of acetylsalicylic and salicylic acid10 can be recycled. All this led to the design of an installation, which worked continuous process and an ecologically-friendly manner. The installation is highly efficient as no product is lost in transferring batches. Furthermore, the recycling of waste reduces the loss of reactants, making it, theoretically, the best method of producing acetylsalicylic acid at the moment.

As said, many other processes are used in manufacturing acetylsalicylic acid. Some employ catalysts, particular solvents or even no solvents at all. What has been demonstrated in this study is that this solvent will not only aid the synthesis phase but also act as a means of separation in the purification phase which is extremely convenient. In the first phase it speeds up reaction (having a similar effect of a catalyst) and is mostly recoverable (none of it ends up in the final product) as it starts evaporating shortly after reaching the temperature needed for the reaction to occur.

In conclusion, the method of synthesis proposed in this paper has been shown to be more efficient than existing methods in terms of efficiency, providing better yields 11 and purer products12 without a separate purification process. The experiments have demonstrated that the idea (process intensification) behind the process is valid and that after necessary refinement, it would be able to be used in an industrial context to manufacture chemicals more cheaply and efficiently.

Acknowledgements

Without the help of many people, the completion of this work would have not been possible:

  • First of all Peter (the editor), and his seemingly endless patience.
  • The LIYSF organization for giving me the opportunity to write for this journal.
  • Prof. Francesca Clerici, for her kind help in clarifying some issues encountered during the work.
  • Maddalena, without whom all of this could have not been possible.
  • My family and many friends that supported me, physically and morally throughout all this work.

To all of them I would like to give my most grateful thanks.