Synthesis of aspirin lab report PDF

Preview

Summary

Download Synthesis of aspirin lab report PDF

Description

The synthesis of aspirin from sodium salicylate and acetic anhydride, followed by analysis through infrared spectroscopy, melting point determination and thin layer chromatography 5PY022 – Medicines In Development and Use

Abstract Aspirin, acetylsalicylic acid, is a non-steroidal anti-inflammatory drug (NSAID) used to treat inflammation, pain, fever and is an important part of the treatment in myocardial infarction (heart attacks). It is produced from the acetylation of salicylic acid, 2-hydroxybenzoic acid, which is derived from salicin, a natural compound found in the bark of willow trees. The purpose of this experiment was to synthesise aspirin from sodium salicylate through the process of acetylation in an esterification reaction using acetic anhydride as the acetylating agent. In order to do this, sodium salicylate was reacted with acetic anhydride and the product was then run through a Büchner funnel using vacuum filtration in order to filter and separate the precipitated product from solution. The synthesised product was then analysed through Thin Layer Chromatography (TLC), resulting in Rf values of both 0.50 (aspirin) and 0.34 (salicylic acid) – as the product had not yet been purified – and Infrared Spectroscopy (IR), highlighting a peak at 1750.96 cm-1. The outcome of this experiment was indeed the production of aspirin, among impurities within the sample. This was confirmed through the similar Rf values between both the synthesised product and pure aspirin obtained from TLC, and the detection of an IR peak belonging to esters from the presence of the C=O ester bond in the IR spectrum of the synthesised product, and thus indicating the presence of aspirin.

Introduction Aspirin is the generic name for acetylsalicylic acid (figure 1). Salicin is aspirin’s precursor compound that is extracted from the bark of the willow tree and is then chemically processed (figure 2) into salicylic acid, 2-hydroxybenzoic acid, which is then converted into aspirin through acetylation (figure 3). Aspirin is a non-steroidal anti-inflammatory drug (NSAID) used to treat a variety of conditions such as fever (antipyretic effects), mild to moderate pain (analgesic effects), rheumatic fever, and acute and long-term inflammatory conditions, such as rheumatoid arthritis, osteoarthritis and pericarditis among others.

Figure 1: acetylsalicylic acid (aspirin) (Sigma-Aldrich, n.d) [1]

Figure 2: the production of salicylic acid from salicin

(Mahdi, 2010)

Figure 3: synthesis of aspirin through esterification of sodium salicylate with acetic anhydride For pain or fever anti-inflammatory, anti-pyretic and analgesic effects typically begin within 30 minutes. (Bartfai and Conti, 2010) indicates that aspirin’s antipyretic effects are a result of it being able to (irreversibly) inhibit the rate-limiting cyclo-oxygenase (COX 1 and COX 2) enzyme in prostaglandin H (PGH) synthesis. This shows that aspirin’s antipyretic effects are indiscriminate of the type of fever-producing substance (pyrogen). Figure 4 shows how aspirin irreversibly inhibits the biosynthesis of prostaglandins (a natural substance which contributes towards inflammation in the body) and cyclic prostanoids such as thromboxane A2 (which contribute towards the clotting of blood) by covalently modifying the enzyme. This is done by acetylating the cyclooxygenase (COX) enzyme’s hydroxyl groups in the variable (R) groups of active site serine amino acid residues at the COX-1 site (Ser530) and the more recently discovered COX-2 site (Ser516) (Ornelas et al., 2017), which less common in the body and is only inducible by growth factors and inflammatory stimuli. It is important to mention that aspirin inhibits COX-1 170x more effectively than it inhibits COX-2, and that aspirin inactivates COX-1 completely when the only change to the COX-2 site is a change in reaction product from PGH2 to 15(R)-HETE (Awtry and Loscalzo, 2000). The inhibitory effect of aspirin on the inducible COX-2 site prevents the biosynthesis of prostaglandins (Ornelas et al., 2017) which reduces the inflammation, consequently resulting in the reduction of symptoms such as swelling, fever, and increased sensitivity to pain which arise as a result of the inflammation.

Figure 4: Aspirin reacts with the amino acid serine’s active site hydroxyl groups in the cyclooxygenase enzyme, rendering it inactive (Alfonso et al., 2014) Aspirin also plays an important part in treatment given shortly after myocardial infarction to reduce the risk of further MI or sudden death, or the risk of stroke in high-risk patients or those who suffer from cardiovascular disease . This is because it has antithrombotic effects and inhibits the cyclooxygenase enzyme through acetylation at the COX-1 site in order to prevent the synthesis of prostanoids such as thromboxane A 2 (a vasoconstrictor and potent platelet activator) (NCGC, 2013) (Giménez-Bastida et al., 2018) from prostaglandin H2 (PGH2) which is derived from the enzymatically catalysed oxidation of arachidonic acid by the prostaglandin-H synthase enzyme (also known as cyclooxygenase). This results in desensitising the function of platelets and so a lowering of the potential to produce blood clots by reducing the aggregation of thrombocytes (Ornelas et al., 2017). Salicylic acid is administered as aspirin – which irritates the stomach to a lesser extent than salicylic acid – and undergoes metabolic biotransformation through hydrolysis back into its active metabolite salicylic acid in the blood, the stomach, small intestine (specifically in the intestinal mucosa) but mostly within the liver when it is taken orally (classing aspirin as a prodrug) and is then absorbed into the bloodstream (Needs and Brooks, 1985). However, aspirin also causes unpleasant side effects to arise as it still causes erosion to the stomach lining, albeit to less of an extent than salicylic acid, and so increases the risk of gastrointestinal bleeding (due to haemorrhaging of the stomach walls) in those taking this medication, especially in those who are on blood-thinning medication such as apixaban, warfarin, heparin, tinzaparin, enoxaparin and rivaroxaban among others (BNF, 2018), take other NSAIDs, are of an older age or drink alcohol. This has therefore sparked the need for the development of enteric-coated formulations with the aim of the tablet dissolving in the small intestine rather than the gastric acid and so a delayed action must be anticipated upon initial use so as to allow time for the tablet to pass through the stomach. Aspirin can be produced from salicylic acid or sodium salicylate by reacting it with a variety of reagents including acetic anhydride or acetyl chloride among others, in an esterification reaction to produce acetylsalicylic acid. These reagents are used as phenols (such as the one within salicylic acid) are much less reactive and much more reluctant to give up a proton than alcohols when subjected to carboxylic acids so react too slowly for preparation purposes, and therefore need a stronger acetylating reagent (aforementioned) in order to initiate a reaction. Sodium acetate or hydrogen chloride gas, respectively, are produced as by-products of each reaction. However, acetic anhydride is preferred as the acetylating agent rather than acetyl chloride because it was the less hazardous reagent. This is because

hydrogen chloride gas, which is toxic if inhaled, is produced upon reaction of acetyl chloride with salicylic acid or any moisture in the air as it is very reactive with water. Thin layer chromatography is very useful for identifying unknown compounds, separating and identifying the components in a mixture, and check how pure a substance is while identifying any impurities within a synthesised product (Touchstone, 1992). This is done by spotting the synthesised product against reference samples of the pure compound that you wish to compare against onto a polar gel plate (typically silica). The organic solvent is allowed to elute up the stationary phase (the silica gel plate) until the solvent front starts to approach the top of the plate (roughly 1cm away) thus carrying the spotted samples up the plate and is then allowed to dry at room temperature. The spots are then viewed under UV light, with the distances they have travelled and the distance the solvent front has travelled up the plate then being marked and recorded. These values can be processed to obtain Rf values which will aid in determining if the crude product is in fact the desired product. The polarity of the compound(s) being tested determines the distance travelled by the samples. Essentially, if the compound is more polar, hydrogen bonds form between polar groups in the compound and polar groups on silica molecules on the surface of the plate thus impeding its elution up the plate. However, the compound travels further up the plate if it is mainly or completely non-polar as there are fewer interactions between polar groups on the surface silica molecules and the compound being analysed.

Methods Molecular weight (g mol-1)

Quantity

Amount (mM)

Sodium salicylate

160.10

2.41 g

15.05

2M NaOH (aq)

40.00

30.0 ml

60.00

Ethanoic anhydride

102.00

5.0 ml

53.00

Aspirin

180.16

2.71 g

15.04

Table 1: amount of chemicals used Table 1 illustrates the different chemicals and the amounts of each used in the experiment. To prepare aspirin, sodium salicylate was reacted with an excess of acetic anhydride at room temperature. A top-pan balance was used to measure 2.41 g of sodium salicylate which was then transferred directly into a weighed 100ml conical flask with 30 ml of 2M aqueous sodium hydroxide, NaOH, then being added in order to convert the sodium salicylate into its even more reactive dianion form. A handful of crushed ice, enough to bring the volume to about 50 ml, was added in to the reaction vessel in order to keep the temperature down. After the ice was added, while continuously stirring the mixture using a magnetic stirrer, 5.0 ml of acetic anhydride was added to the flask and left to mix for 15 minutes. Another handful of crushed ice, enough to bring the volume to about 65 ml, was then added to cool the reaction down while continuously stirring the mixture. With a 3 cm 3 disposable plastic pipette, 15 ml of 4M HCl was then added to the reaction vessel and four further 1 ml

aliquots were added (as precipitation had not occurred yet) so that a heavy white precipitate formed. This was done because any aspirin in the solid product will be encouraged to precipitate faster when strong acid is added due to the fact that the acid acts as the catalyst for this reaction (as the reaction is otherwise slow). This happens because the oxygen on a C=O bond on the acetic anhydride is protonated by the strong acid that is added, increasing the rate of reaction with the salicylic acid as the nucleophilic strength of the phenolic hydroxyl becomes stronger. The acid is classed as the catalyst because the proton returns to the acid at the end of the reaction. The contents of the reaction vessel were then, without delay, run through a Büchner funnel in order to filter and separate the precipitated product from solution using vacuum filtration, and washed with a small volume of cold distilled water to get rid of any residual acetic acid because acetic acid is separated easily from the product upon addition of water as it is very soluble in water. Water was then removed as much as possible from the product, with the product then being placed in a 100 ml beaker which was placed in a drying oven for about 30 minutes and then reweighed in order to estimate the reaction’s yield. In order to analyse the product, various different physicochemical properties of the product were tested in order to determine if these match with that of aspirin’s, as each drug has its own unique physicochemical properties. Thin Layer Chromatography (TLC) (using a commercial silica gel plate) was used to screen the product, the melting point of the product (using a conventional heated block apparatus) was determined, and an IR spectrum of the crude product was produced. To analyse the product using TLC a small quantity of crude product, dissolved in acetone, was spotted onto the plate against solutions of pure aspirin and salicylic acid for comparison and the plate was eluted using a solvent mixture containing 65% hexane, 30% ethyl acetate and 5% acetic acid.

Results Sodium salicylate + acetic anhydride  Aspirin (acetylsalicylic acid) + sodium acetate C7H5O3Na + C4H6O3  C9H8O4 + C2H3O2Na Actual Yield ( g) × 100 Mass of dry crude product: 1.71 g Theoretical yield (g) Theoretical yield: moles salicylate = 2.41 g/160.11 g mol-1 = 0.0151 mol  moles salicylate = moles aspirin as ratio is 1:1 from chemical equation  theoretical mass aspirin = 0.0151 mol Percentage Yield =

Actual Yield × 100 x 180.158 g mol-1 = 2.71 g Theoretical yield % yield: 1.71 g/2.71 g x 100 = 63% Percentage Yield =

Rf of pure aspirin: 1.5 cm/4.4 cm = 0.34 Rf of pure salicylic acid: 2.2 cm/4.4 cm = 0.50 0.32* Rf of crude product: 2.2 cm/4.4 cm = 0.50

Rf =

1.4 cm/4.4 cm = 1.5 cm/4.4 cm = 0.34

distancetravelled by spo di t t ll d b l t

Melting point of pure aspirin: 138 - 140 C (Lewis D., 2003) Melting point range of pure salicylic acid: 158-160 C (Lewis D., 2003) Determined melting points of crude product: 122 oC - 132 oC, 123 oC - 129 oC and 120-133 oC

a)

b)

Figure 5: Thin Layer Chromatography plate a) visualised and b) analysed to obtain Rf values

Figure 6: Infrared spectrum of synthesised crude product

Discussion Through experimentation, a yield of 1.71 g of dry crude product had been obtained. This was much lower than the theoretical yield of 2.71 g that had been calculated, which resulted in the percentage yield being 63%. A low percentage yield was obtained potentially because not all of the sodium salicylate and acetic anhydride starting material had reacted completely as the reaction had not gone to completion so resulted in an amount of product less than that of the theoretical yield. Also, leftover product remaining in the glassware and systematic errors such as the balance not being accurately tared could have also resulted in this reduced yield. Improper weighing of the sodium salicylate may have been caused by not accurately setting the balance at 0.00 g initially as the balance kept fluctuating ±0.01g, meaning more or less sodium salicylate may have been measured than expected. Another reason why “the actual yield is almost always less than the theoretical yield” (Kotz et al., 2009) could potentially be due to loss of product during isolation. What is mean by this is that after separating the crystallised product from solution during the isolation step using

vacuum filtration, not all of the filtered crystals present were adequately removed from the filter paper which may have resulted in a high potential for the product to have remained on the filter paper and be lost. In future, to avoid these sources of errors the initial reading on the balance would be double-checked so as to make sure it is initially at 0.00g precisely, and also ensure there is no product left behind in all glassware and the filter paper while transferring solid during each step of the experiment. The amount of starting materials used could also be increased. Because of the relatively small amount of starting materials used and product, the percentage yield will reflect the significant effect of small losses in solid caused by the many transfers of material. This could be rectified by initially using a greater amount of sodium salicylate and acetic anhydride. The synthesised product’s melting point range was about 16-18 oC lower than that of pure aspirin, and was also greater than 2 oC. The reason for this was because impurities (most likely unreacted salicylic acid, and potentially acetic anhydride) were still present within the sample due to the fact that the synthesised product had not yet been purified through the process of recrystallisation with hot ethanol and so may have caused this characteristic temperature range to have become lowered and broadened. This is because intermolecular bonds in the structure of a solid are disrupted by impurities present in the solid causing them to require less energy to be broken, making the solid easier to melt (Kirsop, 2019). Furthermore, insufficient drying of the product in the drying oven - causing water still present within the product to act as an impurity – could have potentially been another reason for the lowering of the melting point of the product. The process of thin layer chromatography also highlighted the fact that there were impurities within the product. Looking at Figure 5a it can be seen that the spot which corresponds to the synthesised crude product had fragmented into two sections, aspirin being the lower spot and salicylic acid being the upper spot. Figure 5b shows that the distances travelled by the pure aspirin sample (‘A’) was 1.5 cm (yielding an Rf value of 0.34), the pure salicylic acid sample (‘SA’) was 2.2 cm (Rf = 0.50) and 1.4 cm (Rf = 0.32), and the synthesised product (‘P’) was 2.2 cm (Rf = 0.50) and 1.5 cm (Rf = 0.34). The fact that the spot for the product was split into two parts confirms that the product did in fact contain impurities as it had not been purified through recrystallisation with hot ethanol. (Kincaid et al., 1991) instead states that aspirin and salicylic acid have Rf values of 0.67 and 0.70 respectively. The fact that the experimental Rf values obtained were lower than these theoretical values could potentially be because the TLC plate had been left to elute for slightly longer than was needed and so the mobile phase may have travelled up the stationary phase further than was necessary yielding smaller Rf values, and also because impurities may have impeded the progress of the spots on the TLC plate potentially interacting with silica molecules on the plate’s surface. Furthermore, figure 5a also shows that the reference spot corresponding to the pure salicylic acid sample (‘SA’) was split into two parts which could suggest that the sample that had been used had been contaminated with a foreign substance, potentially aspirin, because the spot for 1.4 cm (Rf = 0.32) should not have been present as this correlates with the Rf for aspirin. This could potentially be because some of the pure aspirin sample was mixed in accidentally with the pure salicylic acid sample or due to errors while spotting the sample as

a capillary tube that had already been used to spot aspirin and replaced back in the original container by another student by mistake may have been unknowingly used. In order to avoid this in future, all equipment will be obtained beforehand and keep it clean and keep it separate from other students’ equipment in order to avoid contamination. Another method that could have been employed as part of the experiment to identify if any (unreacted) sodium salicylate impurities were still present is through the use of ferric compounds such as iron (III) chloride, FeCl3, or iron (III) nitrate, Fe(NO 3)3 (both yellow in colour) – the ‘ferric chloride test’ (Ahluwalia and Raghav, 1997). Deprotonated phenols chelate the Fe3+ ligand and, as (Wan et al., 2018) tells us, “phenolics complexed with Fe(III) can be reduced to an Fe(II) complex” thus producing an iron(II)-phenolic complex, [Fe(C6H5O)6]3-. (Ahluwalia and Raghav, 1997) also states that the complex formed is a dark purple/violet colour in solution, which indicates the presence of phenolic groups in the compound as illustrated by the equation 6C 6H5OH + FeCl3  [Fe(C6H5O)6]3- + 3HCl + 3H +. In order to carry out this investigation, the iron (III) chloride could be added to 3 test tubes labelled A, B and C containing pure aspirin and pure salicylic acid solutions (as references) and a sample of the s...