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Using Recrystallisation Improve The Purity Of Aspirin Biology Essay

In this experiment I have investigated the research question – How does the process of recrystallisation improve the purity of Aspirin.

I used a well documented method of preparing Aspirin. Having obtained the Aspirin I performed several recrystallisation processes on it. I then determined accurately the quantity of Aspirin in each of my sample by volumetric analysis. I was then able to determine purity and percentage yield by comparing it to an original tablet Aspirin in the market. I also used melting point to assess purity.

I learnt about Aspirin when we did the chapter medicine and drugs in our class. Aspirin is a very common drug used in our daily lives. The common chemical name is acetylsalicylic acid. Salicylic acid was identified and isolated from the bark of a willow tree but it could not be synthesised in laboratory. In 1893, Felix Hoffman Jr., a chemist found out a practical way for synthesizing an ester derivative of salicylic acid, acetylsalicylic acid. Acetylsalicylic acid, a weaker acid than salicylic acid, was found to have the medicinal properties of salicylic acid without having the objectionable taste or producing the stomach problems as a side effect. The acetyl group effectively masks the acidity of the drug during its ingestion and after it passes into the small intestine, it is converted back to salicylic acid where it can enter the bloodstream and do its pain relieving action [1] . Acetylsalicylic acid is powerful as a pain reliever, fever reducer, and swelling-reducing drug but it also has faults, it causes stomach irritation to some individuals and also may lead to Reye syndrome in young children. I was excited to see if this drug we use has the same purity when prepared in lab and when bought from outside. Aspirin is an important analgesic therefore methods of improving purity are essential. The preparation of Aspirin involves organic synthesis and I learnt about this process when I studied organic chemistry and it an interesting topic to research about. The preparation of drugs that I use in my daily lives excited me particularly as I want to do biochemistry in university and also work in a pharmaceutical company for drug designing. Hence, I decided upon making aspirin in the lab and researching about it.

BACKGROUNG INFORMATION

2.1 Synthesis of Aspirin

The above is the reaction for the formation of Aspirin. This organic synthesis is an esterification reaction between a compound containing a OH group (ester) and an acid. Esters are a type of organic acid in which the hydroxide groups are replaced. The H from the OH group is replaced by a carboxyl carbon ‘C=O’ group. Esterification is the acid catalyzed reaction of a carboxyl (-COOH) group and an -OH group of an alcohol or phenol to form a carboxylate ester. A catalyst is required for the reaction for example concentrated H2SO4.In the synthesis of Aspirin the -OH group is the phenolic -OH group attached to ring of the salicylic acid [2] .

2.2 Purification of Aspirin using the process of recrystallisation

I used the process of recrystallisation to investigate how effective this process is in making Aspirin pure. The process of recrystallisation takes advantage of the relative solubilities of contaminants compared to that of Aspirin [3] . The technique is to use a solvent in which the solid is sparingly soluble at low temperature and quite soluble at higher temperature (at the boiling point of the solvent). In my research Aspirin is insoluble in cold water and hence in the process of recrystallisation I first dissolved Aspirin crystals into hot water and then let it cool down so that it would crystallize out. The solid is dissolved in the minimum quantity of solvent required to produce a solution at the boiling point of the solvent. Upon cooling the solution to room temperature or below, the solid crystallizes out of solution due to its lower solubility at the lower temperature [4] .

Impurities (i.e., any foreign substance) in a solid are classified as soluble or insoluble. The removal of insoluble impurities is accomplished by filtering the hot solution. The insoluble impurities remain on the filter paper. Ideally, soluble impurities remain in solution when the solid being purified crystallizes. (Depending upon concentration and solubility of the impurity in the selected solvent it may sometimes be necessary to recrystallise more than one time. That is some of the soluble impurity might also crystallize. If any soluble impurity crystallizes, the melting point of your product will be depressed). When recrystallisation is complete the purified solid is isolated by filtration and the crystals are washed with a small quantity of cold solvent (to rinse off the solution of soluble impurities coating the freshly filtered solid) [5] .

2.3 Determination of purity using melting point apparatus

I also used melting point determination to give further evidence towards the purity of aspirin. Melting point is a useful measure for the purity of a solid. “Melting point apparatus” is commonly used for this purpose. It consists of a heated metal block with holes for a thermometer and melting point tubes. The capillary tubes are provided open-ended and the crystalline solid can be transferred into the tube and forced to the bottom with gentle tapping. The compound is heated slowly especially around its melting point for accuracy.

There are attractive forces (intermolecular interactions) between the molecules in a solid that keep them together in an ordered crystalline structure. If enough heat energy is added to the solid the internal kinetic energy of the molecules causes them to move in the solid. At the temperature where the energy of molecular motion overcomes the attractive forces between molecules the compound begins to melt. When a solid is pure the molecules are all identical and thus the interactions between molecules are similar and thus the sample will melt at a distinct temperature. Impure compounds, on the other hand, have a range of intermolecular interactions between molecules and will melt over a range of temperatures. [6] 

3. APPARATUS [7] :-

3.1 Equipments:

The apparatus listed below does not list quantities for repeat readings.

Conical flask (100 cm3) (×1)

Measuring cylinders (10 cm3) (±0.5cm3) (×2)

Beaker (100 cm3) (×2)

Glass rod (×1)

Vacuum filtration flask (×1)

Rubber tubing for vacuum flask (×1)

Hirsch funnel (×1)

Water bath containing crushed ice (×1)

Source of hot water (×1)

Test-tubes (×4)

Meltemp apparatus – for finding the melting point of Aspirin

Burette (50cm3) (×1)

Clamp stand (×1)

Spatula (×2)

Watch glass (×1)

Melting point capillary tube (×1)

Filter paper to fit Hirsch funnel (×1)

3.2 Chemicals:

2-hydroxybenzoic acid (salicylic acid) (2g)

Ethanoic anhydride (4cm3)

Concentrated sulphuric acid (5 drops)

Ethanoic acid (glacial) (4cm3)

(1)Aspirin tablet

Phenolphthalein solution

Sodium hydroxide solution (0.1 mol dm-3)

95% alcohol

4. DIAGRAM:

4.1 Hirsch Funnel:

4.2 Melting Point Apparatus:

5. METHOD [8] :-

‘‘Shake 2g of 2-hydroxybenzoic acid (salicylic acid) (CARE Irritant) with 4 cm3 of ethanoic anhydride (CARE Corrosive) in a 100 cm3 conical flask.

Add 5 drops of concentrated sulphuric acid (CARE Corrosive) and continue agitating the flask for about 10 minutes. Crystals of Aspirin will appear and soon the whole will form a crystalline mush.

Dilute by stirring in 4cm3 of cold glacial ethanoic acid (CARE Corrosive) and cool by placing in a water bath containing crushed ice.

Filter off the crystals using a Hirsch funnel (a small funnel for vacuum filtration), washing once with ice cold water to remove residual acid.

Place the crude Aspirin in a 100cm3 beaker. Add hot, but not boiling, water until it dissolves. A mass of very pure Aspirin crystals will form; cool the flask by surrounding it with cold water’’.

Filter them again and rinse the crystals with the chilled water.

‘‘The insoluble impurities remain on the filter paper and the filtrate contains the product. Aspirin can be recovered from this solution by evaporation of the recrystallisation. [9] ’’ Leave the crystals overnight on a watch glass to dry completely. This process is known as recrystallisation and is a way of purifying a solid product (Aspirin).

Do the recrystallisation process three times and after every recrystallisation remove some sample of Aspirin and store in a test tube to test later.

Now do titration of the samples stored after each recrystallisation. Take some of Aspirin for each sample and leave some in the test tube for testing the melting point.

For the process of titration, take the Aspirin from each sample into a 50 cm3 conical flask and dissolve it in 5 cm3 of 95% alcohol and add two drops of phenolphthalein solution to it.

Titrate the solution in the conical flask with 0.1 mol dm-3 sodium hydroxide from a burette (CARE Eye protection must be worn).

Record the volume needed to produce the first tinge of pale pink colour in the indicator. This measure the end-point of the titration.

‘‘Take a capillary tube 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.

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. [10] ’’

Then do the titration of an original tablet of Aspirin available in the market. Then test the melting point of the original tablet of Aspirin by the method described above.

Compare the melting point which you get from the samples and the original tablet of Aspirin with the one given in the data booklet.

6. OBSERVATIONS:-

When I mixed salicylic acid with ethanoic anhydride, the solution turned milky. When to the solution I added concentrated sulphuric acid, the solution turns colourless and then after agitating for 10 minutes the solution again turns milky white. The beaker is hot and hence we can say that the reaction between concentrated sulphuric acid and the solution (ethanoic anhydride + salicylic acid) is exothermic.

When I was doing my melting point I saw that the solid obtained after the first recrystallisation actually turned black before actually getting close to the melting point of the original Aspirin. As the number of recrystallisation increased I could see that the melted Aspirin was still white and was getting closer to the melting point of the original Aspirin (135°C) as mentioned in the data book.

7. DATA COLLECTION AND PROCESSING:-

7.1 The data of titrations of different recrystallisation samples of Aspirin:-

7.1.1 Original Aspirin tablet

Burette solution (cm3)

0.1 mol dm-3 sodium hydroxide solution

Indicator

Phenolphthalein solution

Trial

1st reading

2nd reading

3rd reading

Burette readings (cm3)

Final (±0.1)

44.5

44.0

44.0

44.3

Initial (±0.1)

69.0

69.0

69.0

69.0

Volume used (titre) cm3 (±0.2)

25.0

25.0

I have not used as they are not concordant.

I have used these reading for my mean titre.

Mean titre (cm3) (±0.2)

25.0 + 25.0 = 50.0

50.0 ÷ 2 = 25.0 (mean titre)

Volumetric calculations

Volume of NaOH used = 25.0 cm3.

Moles of NaOH n = CV V = 25.0 cm3 = 25.0 ÷ 1000 = 0.025 dm3 n = 0.1 × 0.025 = 0.0025 mol

So, moles of Aspirin will also be equal to 0.0025mol because the reaction ratio between NaOH and Aspirin is 1:1.

Weighed out sample of Aspirin = 0.62 g

How many grams of Aspirin reacted with NaOH?

Aspirin = C9H8O4

g = n × Mr

= 0.0025 × Mr [(12.01 × 9) + (1.01 × 8) + (16.00 × 4)]

= 0.0025 × 180.17

= 0.45 g

Percentage of Aspirin reacted = (0.45 ÷ 0.62) × 100

= 73%

7.1.2 Aspirin after 1st recrystallisation

Burette solution (cm3)

0.1 mol dm-3 sodium hydroxide solution

Indicator

Phenolphthalein solution

Trial

1st reading

2nd reading

3rd reading

Burette readings (cm3)

Final (±0.1)

3.4

3.6

3.7

3.8

Initial (±0.1)

0.0

0.0

0.0

0.0

Volume used (titre) cm3 (±0.2)

3.6

3.7

3.8

I have used these reading for my mean titre.

Mean titre (cm3) (±0.2)

3.6 + 3.7 + 3.8 = 11.1

11.1 ÷ 3 = 3.7 (mean titre)

Volumetric calculations

Volume of NaOH used = 3.7 cm3.

Moles of NaOH n = CV V = 3.7 cm3 = 3.7 ÷ 1000 = 0.0037 dm3 n = 0.1 × 0.0037 = 0.00037 mol

So, moles of Aspirin will also be equal to 0.00037mol because the reaction ratio between NaOH and Aspirin is 1:1.

Weighed out sample of Aspirin = 0.30 g

How many grams of Aspirin reacted with NaOH?

Aspirin = C9H8O4

g = n × Mr

= 0.00037 × Mr [(12.01 × 9) + (1.01 × 8) + (16.00 × 4)]

= 0.00037 × 180.17

= 0.066 g

Percentage of Aspirin reacted = (0. 066 ÷ 0.30) × 100

= 22%

7.1.3 Aspirin after 2nd recrystallisation

Burette solution (cm3)

0.1 mol dm-3 sodium hydroxide solution

Indicator

Phenolphthalein solution

Trial

1st reading

2nd reading

3rd reading

Burette readings (cm3)

Final (±0.1)

12.6

12.1

12.1

12.3

Initial (±0.1)

9.2

9.2

9.2

9.2

Volume used (titre) cm3 (±0.2)

2.9

2.9

I have not used as they are not concordant.

I have used these reading for my mean titre.

Mean titre (cm3) (±0.2)

2.9 + 2.9 = 5.8

5.8 ÷ 2 = 2.9 (mean titre)

Volumetric calculations

Volume of NaOH used = 2.9 cm3.

Moles of NaOH n = CV V = 2.9 cm3 = 2.9 ÷ 1000 = 0.0029 dm3 n = 0.1 × 0.0029 = 0.00029 mol

So, moles of Aspirin will also be equal to 0.00029mol because the reaction ratio between NaOH and Aspirin is 1:1.

Weighed out sample of Aspirin = 0.15 g

How many grams of Aspirin reacted with NaOH?

Aspirin = C9H8O4

g = n × Mr

= 0.00029 × Mr [(12.01 × 9) + (1.01 × 8) + (16.00 × 4)]

= 0.00029 × 180.17

= 0.052 g

Percentage of Aspirin reacted = (0. 052 ÷ 0.15) × 100

= 35%

7.1.4 Aspirin after 3rd recrystallisation

Burette solution (cm3)

0.1 mol dm-3 sodium hydroxide solution

Indicator

Phenolphthalein solution

Trial

1st reading

2nd reading

3rd reading

Burette readings (cm3)

Final (±0.1)

17.2

17.6

17.9

17.9

Initial (±0.1)

13.1

13.1

13.1

13.1

Volume used (titre) cm3 (±0.2)

I have not used as they are not concordant.

4.8

4.8

I have used these reading for my mean titre.

Mean titre (cm3) (±0.2)

4.8 + 4.8 = 9.6

9.6 ÷ 2 = 4.8 (mean titre)

Volumetric calculations

Volume of NaOH used = 4.8 cm3.

Moles of NaOH n = CV V = 4.8 cm3 = 4.8 ÷ 1000 = 0.0048 dm3 n = 0.1 × 0.0048 = 0.00048 mol

So, moles of Aspirin will also be equal to 0.00048mol because the reaction ratio between NaOH and Aspirin is 1:1.

Weighed out sample of Aspirin = 0.15 g

How many grams of Aspirin reacted with NaOH?

Aspirin = C9H8O4

g = n × Mr

= 0.00048 × Mr [(12.01 × 9) + (1.01 × 8) + (16.00 × 4)]

= 0.00048 × 180.17

= 0.086 g

Percentage of Aspirin reacted = (0. 086 ÷ 0.15) × 100

= 57%

7.1.5 Aspirin after 4th recrystallisation

Burette solution (cm3)

0.1 mol dm-3 sodium hydroxide solution

Indicator

Phenolphthalein solution

Trial

1st reading

2nd reading

3rd reading

Burette readings (cm3)

Final (±0.1)

21.3

20.9

21.0

21.1

Initial (±0.1)

16.9

16.9

16.9

16.9

Volume used (titre) cm3 (±0.2)

4.0

4.1

4.2

I have used these reading for my mean titre.

Mean titre (cm3) (±0.2)

4.0 + 4.1 + 4.2 = 12.3

12.3 ÷ 3 = 4.1 (mean titre)

Volumetric calculations

Volume of NaOH used = 4.1 cm3.

Moles of NaOH n = CV V = 4.1 cm3 = 4.1 ÷ 1000 = 0.0041 dm3 n = 0.1 × 0.0041 = 0.00041 mol So, moles of Aspirin will also be equal to 0.00041mol because the reaction ratio between NaOH and Aspirin is 1:1.

Weighed out sample of Aspirin = 0.10 g

How many grams of Aspirin reacted with NaOH?

Aspirin = C9H8O4 g = n × Mr

= 0.00041 × Mr [(12.01 × 9) + (1.01 × 8) + (16.00 × 4)]

= 0.00041 × 180.17

= 0.074 g Percentage of Aspirin reacted = (0. 074 ÷ 0.10) × 100 = 74%

7.2 The data of melting points of different recrystallisation samples of Aspirin:-

7.2.1 Melting point after first recrystallisation

Number of recrystallisation

Temperature (ºC) (± 0.1ºC)

Original melting point of Aspirin (ºC)

Trial

1st reading

2nd reading

3rd reading

Average

1

155.0

152.5

151.9

151.7

152.0

135.0

2

154.1

150.0

148.5

148.2

148.9

135.0

3

115.2

120.9

122.8

122.9

122.2

135.0

4

124.7

125.4

126.1

126.9

126.1

135.0

Original tablet

128.0

128.8

129.2

129.9

129.3

135.0

7.2.2 Graph showing the difference between melting points of Aspirin which was prepared and recrystallised in lab and melting of Aspirin from the data book

Y-axis = temperature (in ⁰C)

X-axis = number of recrystallisations of aspirin samples prepared in lab and aspirin available in market

8. INTERPRETATION OF THE DATA

I will now explain the results of melting point and titration. From the results of titrations we can see a trend flowing and how after each recrystallisation the sample gets purer. As the quantity of Aspirin decreased after each recrystallisation, the mass of Aspirin in that quantity was more compared to the previous recrystallisation. When compared to the original tablet which we get in the market, I could speculate that the producer has done almost four recrystallisations to get that purity of Aspirin. More pure Aspirin can be obtained if more recrystallisations are done.

We could say that the difference in melting point might be higher because it might contain impurities like unreacted salicylic acid or other by-products of the reaction or decomposition products. We can see that the difference in the melting point is getting less as the number of recrystallisation increase and closer to the melting point of Aspirin published in the data booklet. Further evidence to my theory that the number of recrystallisations increases the purity is my melting point data. From my graph it can be seen clearly that as the number of recrystallisations increase the closer to the melting point of pure aspirin we get in the market.

My data shows that the percentage purity of aspirin increased with each recrystallisation process. For example, after the first recrystallisation the percentage purity was only 22%, however when fourth recrystallisation was done the percentage purity was 74% showing a significant increase. This can be seen in the following graph:-

9. CONCLUSION

In answer to my research question, ‘How does the process of recrystallisation improve the purity of Aspirin?’ I have found significantly that the percentage purity increases with each recrystallisation and this is evident in my graph under the heading ‘‘Interpretation of the data.’’

Recrystallisation is an important technique in organic Chemistry. The general method is to find a solvent that dissolves the product more readily at high temperature than at low temperature, make a hot solution, and allow to crystallise on cooling. The crude product might contain; impurities which are insoluble in the solvent; impurities which are slightly soluble in the solvent; and impurities which dissolve readily in the solvent. The solvent itself has also to be removed or it behaves as an impurity in its own right. It must not leave behind any residue. One simple way to tell whether an organic compound is pure is to measure its melting (or boiling) point. A pure compound melts sharply: if impurities are present it melts slowly (over a range of temperature). [11] 

The process of recrystallisation in my experiment increased the purity of Aspirin but with a decrease in the quantity produced. The solid will readily dissolve in a larger quantity of solvent; the larger the volume of solvent the greater the loss of product [12] . This is the reason why after every recrystallisation I lose Aspirin. The process of recrystallisation removes the impurities present and this can be concluded from the fact that the difference between the melting point of the sample and the melting point of Aspirin from data booklet decreases with each recrystallisation.

I could conclude that a pharmaceutical company should always have a balance between producing a very pure product, which means many recrystallisation processes and producing enough quantity of the product to make it a profitable industry.

10. EVALUATION

10.1 Random error

The apparatus I used had uncertainties like the measuring flask has an uncertainty of ±0.5cm3, the burette had an uncertainty of ±0.1cm3 and the melting point apparatus also had an uncertainty of ± 0.1ºC. This results in errors in my results. The equipment error could be reduced by using equipments with less error. For example, I could use a burette with an error of ±0.05 cm3 instead of a burette with error of ±0.1 cm3.

10.2 Systematic error

In the process of titration there could a parallax error caused if the reading from the burette is not read at eye-level. So, when taking the reading from the burette, the level of the eye should be same as the level of the meniscus. When we are titrating different samples of recrystallisations the colour of the indicator changes from pale pink to dark and it is difficult to know the end-point of the titration process.

10.3 Modifications in the method of preparation of Aspirin and its recrystallisation

I modified the method to improve it in the following ways: - After each recrystallisation there was loss of Aspirin and the decrease in the mass of Aspirin limited the number of recrystallisations needed to get the most pure form of Aspirin. So, if I doubled the mass of reactants I can get doubled the mass of Aspirin produced and an increase in the number of recrystallisations. When I mixed salicylic acid and ethanoic anhydride solution in concentrated sulphuric acid it is hard to get the formation of a crystalline mush of Aspirin by agitating the flask. Hence, instead I used a magnetic stirrer which gives a uniform stirring and all the chemicals are mixed properly. I found this to be more effective at producing the mush. During the process of recrystallisation a lot of Aspirin is lost. When I used the filter paper, I cut it the same size as the Hirsh funnel. Instead I could have used a bigger filter paper in the funnel so that the impurities do not leak out of the edges of the filter paper and I could get a purer sample of Aspirin. When the Aspirin is left overnight to dry in an evaporating dish, there could be many contaminants which would get mixed in the Aspirin and if this is not taken into consideration in pharmaceutical industries then this could lead to serious health problems. Hence, the Aspiring should be covered when left overnight to dry. I could use different methods of purification of Aspirin like thin layer chromatography which is a sensitive and quick way of detecting impurities in an organic product (Aspirin). I could also use spectroscopy which provides a very good method for analyzing an organic compound. By comparing the infra-red spectrum for Aspirin with the spectrum of compound in a database I can check on its purity.

11. UNANSWERED QUESTIONS

Unfortunately, not all my questions could be answered in this experiment. It would have been interesting in further researching and comparing more brands of Aspirin available in the market, if the producers just recrystallised the sample of Aspirin twice so that they do not lose a lot of their product in the process of recrystallisation or the producers actually tried to produce a pure sample of Aspirin ignoring the decrease in yield after each recrystallisation and considering the fact that this can affect the health of humans.

When an organic compound has been made it needs to be purified, particularly if it is a pharmaceutical chemical. This is because most organic reactions produce by-products but, even if the reaction is a ‘clean’ one, the purity standards for many products are so stringent that small amounts of other compounds have to be removed. In particular the catalyst used in this reaction is concentrated sulphuric acid and must all be removed.

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