Testing The Purity Of Aspirin Biology Essay

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Aspirin, or acetylsalicylic acid, is a drug used to relieve minor aches and pains or fever. It originates back to the time of Hippocrates who discovered that chewing the bark of the willow tree helped relieve pain. Later on in 1829, scientists discovered that it was in fact the compound salicin (see below) which gave you pain relief. Then in 1838 Rafael Piria (an Italian scientist) split salicin into a sugar and aromatic component called Salicylaldehyde (see below). Then, by hydrolysis and oxidation he converted salicylaldehyde into an acid of crystallised needles which he named salicylic acid (see below) [2]. Salicylic acid is one of the main chemicals needed to synthesise aspirin. However, salicylic acid was tough on the lining of the stomach so scientists had to find a way to 'buffer' the compound. In 1853 Charles Gerhardt neutralised salicylic acid by reacting it with sodium and acetyl chloride creating acetylsalicylic acid which we now know as aspirin (see below):

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Synthesising aspirin directly from the bark of the Willow tree would have proven extremely time consuming and impractical. It also would have no bearing on any of the aims which I have set out for my investigation whether I synthesised aspirin directly from Willow bark, or started at a different point in the process. As a result of my research into the origins of aspirin, I found that aspirin can be synthesised from Salicylic acid which is readily available within the school. I decided to start with Salicylic acid and synthesise my aspirin from there.

Aromatic compounds

Salicylaldehyde is part of the Aromatic compounds which are sometimes known as Arenes. They are unsaturated and contain a benzene ring. The aldehyde part of salicylaldehyde comes from the fact that it has a double bonded oxygen and a singly bonded hydrogen attached and an alcohol group bonded to a carbon atom. Salicylaldehyde is used in the production of salicylic acid.

[Diagram of Salicylaldehyde with benzene ring circled in red and aldehyde group circled in blue]

Hydrolysis and oxidation

Hydrolysis is the addition of water to split a compound. For example, CH3CH2COOCH2CH3 è CH3CH2COOH + CH3CH2OH. Water has been added and the ester bond has been broken. Oxidation on the other hand is the loss of electrons from a compound. This is indicated by a change in the charge of an atom. For example the oxidation of Sodium can be written as Na è Na+ + e-. The charge on the Na atom has gone from 0 to +1, therefore the oxidation state has increased by one and the atom has been oxidised. Salicylic acid is produced from Salicylaldehyde by hydrolysis and oxidation.

Nucleophilic Substitution

One way in which aspirin can be synthesised from salicylic acid is through Nucleophilic substitution. The salicylic acid reacts with ethanoic anhydride to form an ester [6]:

C7H6O3 + C4H6O3 è C9H8O4 + C2H4O2

Salicylic acid + Ethanoic anhydride è Aspirin + Ethanoic acid

A nucleophile is "a molecule or negatively charged ion with a lone pair of electrons that it can donate to a positively charged atom to form a dative bond." This is true in the case of the salicylic acid molecule [7]. The way in which the substitution is carried out is detailed below.

The oxygen atoms on the ethanoic anhydride are more electronegative than carbon. As a result, the oxygen atoms attract electrons more strongly and become δ- while the carbon atom is δ+. The oxygen atom from the phenol is also δ- and the lone pair of electrons is attracted to the δ+ carbon atom. The salicylic acid acts as a nucleophile and the lone pair on the oxygen atom attacks the δ+ carbon atom. This results in the substitution of a hydrogen atom from the phenol group, with the CH3CO- group of the ethanoic anhydride. The remaining products are Aspirin and ethanoic acid [5].

[5] (old.iupac.org/publications/cd/medicinal_chemistry/Practica-I-11.pdf)

[6] (www.alevelchem.com)

[7] (Chemical ideas 13.5)

Aim 2a: Purify the Aspirin

Purifying my synthesised sample of Aspirin is essential as it is very likely to be impure. This could be due to solid objects from around the lab, or filter paper from drying. Also, contamination of pipettes could lead to unwanted chemicals being incorporated into the sample. There may also be unreacted salicylic acid, or some aspirin may have become hydrolysed and formed salicylic acid (see below).

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The method I used to purify my sample of aspirin was recrystallisation. This is because it is the method used for purifying organic solids. It is also not very time consuming and is relatively easy to do. It yields very good results as recrystallisation serves the purpose of removing insoluble impurities as well as soluble impurities.

Aim 2b: Identify that the product formed is Aspirin

It is important to ensure that the purified product is actually Aspirin. This is because if it is not Aspirin, then testing the purity of it will become useless. It is possible for the Aspirin to be converted back into salicylic acid due to hydrolysis.

Hydrolysis of Aspirin

Hydrolysis is the addition of water to split a compound. In the case of aspirin:

C9­H8O4 + H2O  C7H6O3 + CH3COOH

I had to heat the water before adding aspirin to dissolve the aspirin without hydrolysing it. I also had to make sure that I didn't heat the aspirin for too long, or let the temperature get too hot. This made sure that the aspirin didn't hydrolyse back into salicylic acid.

When water is added to aspirin, it splits the ester bond that is formed when salicylic acid reacts with ethanoic anhydride. However, since the ethanoic anhydride splits when it first reacts with aspirin, salicylic acid and ethanoic acid are formed, rather than salicylic acid and ethanoic anhydride.

Diagram showing hydrolysis of aspirin with ester bond circled:

Aim 3: Test the purity of the purified aspirin

Testing the purity of my sample of aspirin is an indicator of how well the recrystalisation method if purification worked. To test my sample of aspirin, I made up standard solutions and titrated my aspirin against pure aspirin and aspirin tablets from the supermarket. This is called an assay and would give me a good indication of how pure my aspirin was compared to other sources of aspirin. Another method I used was Thin-Layer Chromatography (T.L.C.). This allowed me to distinguish how close my aspirin was to pure aspirin and also how close it was to salicylic acid. I did this by measuring R­f values. I also sent my sample of Aspirin to GSK for NMR (Nuclear Magnetic Resonance) spectroscopy.

Method: Synthesise Aspirin

Equipment

Fume cupboard

150cm3 conical flask

20cm3 measuring cylinders

Apparatus for vacuum filtration

Buchner funnel

Salicylic acid

Ethanoic anhydride

Concentrated Sulfuric(VI) acid

Glacial Ethanoic acid

Water bath containing crushed ice

Funnel

Method

Working in a fume cupboard, I swirled 10.0534g of Salicylic acid with 20cm3 ethanoic anhydride in a 150cm3 conical flask.

I then added 25 drops of concentrated sulfuric(VI) acid (to act as a catalyst). I then continued to agitate the flask until crystals of aspirin appeared, forming a 'crystalline mush.'

At that point, I stirred in 20cm3 of cold glacial ethanoic acid. This diluted the mixture. I then placed the solution into a bath of crushed ice and water.

I then filtered off the crystals using a Hirsch funnel and vacuum filtration.

Afterwards I dried the sample on filter paper in an oven for 5 minutes. I then measured the sample and obtained 11.0062g of aspirin.

[Diagram of Vacuum Filtration]:

[8] (http://en.wikipedia.org/wiki/File:Vacuum-filtration-diagram.png)

Percentage Yield

The percentage yield will help determine the effectiveness of this method for synthesising aspirin. To calculate percentage yield, we have to start with the equation for the synthesis of Aspirin from Salicylic acid:

C7H6O3 + C4H6O3 è C9H8O­4 ­+ C2H4O2

(Salicylic acid) (Aspirin)

During my investigation, I used 10.0534g of Salicylic acid and obtained 11.0062g of Aspirin. The Relative Molecular Mass of Salicylic acid is 138 and the Relative Molecular Mass of Aspirin is 180.

Therefore, the theoretical maximum yield is (180/138)*(10.0534) = 13.1131g.

The actual yield was 11.0062. From this information we can calculate the percentage yield which is equal to the actual yield/maximum theoretical yield multiplied by 100 to give a percentage.

From the results which I obtained, I calculated the percentage yield to be (11.0062/13.1131)*100 = 83.9%

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I can conclude, based on my results, that this is an effective method of synthesising aspirin.

[9] (WM 5.1)

Methods: Purify the Aspirin

Equipment

My sample of aspirin

A suitable solvent (H2O)

Filter paper

Conical flask

Buchner funnel

Vacuum filtration equipment

Funnel

Beaker

Bunsen burner

Heatproof mat

Gauze

Tripod

Glass Rod

Method

Using the Bunsen burner and equipment, I heated the minimum amount of water it would take to dissolve my aspirin. I then dissolved my sample of aspirin in the hot water.

I then poured the hot solution into a conical flask through filter paper in a funnel. This filtered out any insoluble impurities.

Afterwards, I left the solution in the conical flask to cool. This allowed solid crystals to form.

I then used vacuum filtration to filter out any soluble impurities, washing once with cold water.

Finally, I scraped off the crystals from the top of the Buchner funnel and left on filter paper, and put in an oven for 15 minutes to dry.

The reason that this method works is because the aspirin will dissolve in the water due to hydrogen bonding (see below), however, solid impurities will not dissolve. When the solution is filtered, the solution will pass through but any solid will not, this means that solid impurities will be removed.

The soluble impurities are also removed. The reasoning is that soluble impurities will remain in solution while the aspirin crystallises out. As a result, the soluble impurities will filter through during vacuum filtration whereas the insoluble aspirin will remain.

Hydrogen Bonding Properties of Aspirin

Hydrogen bonding occurs when there is a large dipole between a Hydrogen atom and Fluorine, Oxygen or Nitrogen atom which it is bonded to. The result is that the Hydrogen is slightly positively charged (δ+) whereas the atom that it is bonded to is slightly negatively charged (δ-). The dipoles arise due to the large electronegativity values of Fluorine, Oxygen, and Nitrogen. Electronegativity is the measure of how strongly an atom can attract electrons. In Hydrogen Fluoride, Fluorine, for example, will attract electrons more strongly than Hydrogen and a dipole will be formed as a result.

Diagram of Fluorine attracting electrons more strongly than Hydrogen and a dipole being formed.

δ+ δ-

This dipole can be seen in the COOH group of aspirin. The oxygen atom is more electronegative than the hydrogen atom so a dipole is formed.

As well as being δ-, the oxygen atom has a lone pair of electrons which allows hydrogen bonding to occur. The δ+ hydrogen atom is attracted to the δ- oxygen atom and the lone pair of electrons on the oxygen atom forms the bond with the hydrogen atom.

Diagram of Hydrogen Bonding of Aspirin and a Water molecule:

[10] (CI3.1)

Methods: Identify that the product formed is Aspirin

Mass Spectrometry

Mass spectrometry is used to measure the molecular mass of different particles, such as atoms or molecules, in a sample. The results can then be analysed and a structure for the molecule can be predicted. A mass spectrometer has five main components:

Sample Inlet: This is where the sample to be analysed is inserted into the mass spectrometer.

Ionisation Area: The sample to be analysed must be ionised first to form a positively charged ion. This is done by a stream of high-energy electrons bombarding a sample. This knocks electrons off the outside of the atoms in the sample and produces positive ions:

X(g) + e- è X+(g) + 2e-

The atoms in the sample must be ionised as an electric field accelerates them in the 'acceleration area.' This can only be done with charged particles. The ions are produced and accelerated as separate pulses. Also, the sample must be charged in order to be detected by the ion detector.

Acceleration Area: The elements are accelerated by an electric field of known strength, accelerating all particles to the same level of kinetic energy. This is important as changes in mass can be detected when all the particles have the same kinetic energy.

Drift Region: All the air molecules in the drift region have been pumped out creating a strong vacuum. This allows the ions to move along the region without hitting anything else. The reason that this is important is because the velocity of the particles is measured to give an idea of the mass. The equation for kinetic energy is:

Ek = ½ mv2

Where 'm' is the mass of the ion and 'v' is the velocity. If we know the kinetic energy and the velocity, we can predict the mass be rearranging the formula to:

m = (2Ek)/v2

The kinetic energy is known as the electric field accelerates the particles to a known kinetic energy value. The velocity is detected by the ion detector. Therefore, since m = (2Ek)/v2, and the kinetic energy is constant, larger molecules will have a lower value for velocity. This means that larger molecules will take longer to move through the drift region.

Ion Detector: After they pass through the drift region, the various ions are detected by an ion detector. This works because an electrons flowing in the detector become attracted to the previously ionised atoms. This produces a varying electric current. The detector calculates the mass from the information of kinetic energy and the time taken for the ions to reach it. It then converts the information into a mass spectrum which shows the masses of the ionised molecules and their intensity.

[Diagram of a time-of-flight mass spectrometer]

[11]

A mass spectrum is shown on a graph of m/z against percentage intensity. The m in the ratio is mass while the z means charge. Since the charge=+1, the m/z ratio is equivalent to the mass of the ion. Each peak corresponds to a mass and the intensity that at which it is present in the sample. The peak furthest to the right (the one with the largest mass) corresponds to the full sample with just one electron removed. It is also known as the M+ peak. The other peaks correspond to sections of the full sample with one electron removed. These are called fragments. This means that we can predict the structure of the sample by working backwards from the fragments and by taking into account the largest mass. For example, consider the mass spectrum of pentane:

[13]

The Relative Molecular Mass (RMM) of pentane - which is found by adding the RMM of all the atoms that make it up - is 72 (it is a relative scale so there are no units). This corresponds with the peak which is the furthest to the right with a m/z of 72. However, there are also peaks with lower m/z ratios. These are results of fragmentation. For example, the peak with an m/z of 57 has a mass of 15 less than the peak at 72. This could mean that a CH3 molecule has been lost from the molecule with m/z of 72. We can then apply this practice to the other peaks and work out which molecules have been lost from the main peak. Therefore, working backwards through fragmentation, we can predict the structure of the M+ molecule [11,12].

[11] (CI 6.5)

[12] (http://www.chemguide.co.uk/analysis/masspec/howitworks.html)

[13] (http://www.chemguide.co.uk/analysis/masspec/fragment.html)

Melting Point Equipment List

Recrystallised Aspirin

Pure Aspirin

Salicylic acid

Capillary tubes

Thermometer

Melting Point Apparatus (Electronic)

Bunsen burner

Heatproof mat

Melting Point Method

Using the Bunsen burner, I melted the ends of 3 capillary tubes. This ensured that when the capillary tube was inserted into the melting point apparatus, the sample did not fall out of the end.

I then placed some of my recrystallised aspirin into the open end of the capillary tube and flicked it so that it would fall to the bottom where the end was closed. I then repeated this with salicylic acid and pure aspirin.

I then place each of the capillary tubes and a thermometer into the melting point apparatus and waited for the temperature to rise.

I recorded the points at which each sample began and finished melting. I compared this melting range to published values for the melting points of the substances.

Melting point of Aspirin

Aspirin has the structure:

The Carboxylic acid (COOH) group on the Aspirin allows molecules of aspirin to form hydrogen bonds with each other (see hydrogen bonding in aspirin section). This means that the melting point of aspirin will be relatively high. Permanent dipole - permanent dipole bonds may also be formed from the C=O bonds on the aspirin molecule. These bonds occur when one atom is more electronegative than another and a dipole is set up. Opposite charges attract which causes a bond to form.

Salicylic acid, on the other hand, has the structure:

There are now two opportunities for hydrogen bonding from the COOH group and the COH group. As well as this there is an opportunity for permanent dipole - permanent dipole bonds from the C=O bonds. As a result, the melting point of salicylic acid is expected to be higher than the melting point of aspirin as there are more hydrogen bonds formed which means that more energy is needed to break the bonds.

[14] (CI 5.3)

[15] (CI 5.4)

Methods: Test the Purity of the Purified Aspirin

Originally, I used approximately 0.5g of each sample for my titrations. However, this led to me gaining very small titres. As a result, a small change in measurements led to a large change in percentage purity. Also, the percentage errors for measurements that small would be extremely large. This means that the accuracy of my results would be compromised massively. Therefore, I changed my method and used a larger amount of each sample to titrate with. The final method is outlined below:

Assay Equipment List

Mortar and Pestle

Burette

100cm3 conical flasks

10cm3 measuring cylinders

Aspirin tablets

Recrystallised Aspirin

Pure Aspirin

95% Ethanol

Sodium Hydroxide solution, 0.1 mol dm-3

Phenolphthalein indicator

Assay Method

I ground up some supermarket Aspirin and weighed 1.462g and transferred it to a 100cm3 conical flask. I then weighed 1.572g of my sample of aspirin and transferred it to a different 100cm3 conical flask. Finally I weighed 1.477g of pure aspirin and transferred it to a third 100cm3 conical flask.

I used 50cm3 volumetric flasks to measure out 50cm3 of 95% ethanol and then added it to each of the conical flasks. I then swirled the flasks until all the aspirin had dissolved to make up three standard solutions.

I then transferred 10cm3 of each solution into a different conical flask.

Afterwards, I added a few drops of phenolphthalein indicator to each.

I then placed each conical flask in turn underneath the burette. I allowed 0.1 moldm-3 sodium hydroxide to flow from the burette into each conical flask and recorded the volume needed to produce a permanent pink colour for over 10 seconds.

Finally, I repeated the test 3 more times with each solution and found an average titre for each.

[16] (WM 6)

Why is a Pink colour formed?

The pink colour that forms is an essential part of the assay. It is used to help determine the purity of the aspirin by determining how much aspirin is in the sample. The reason it works is because aspirin is an acid. Aspirin is an acid because of the carboxylic acid group which gives the molecule a relatively stable anion (negatively charged form) which means that it can readily donate a proton. The Bronstead-Lowry theory states that "an acid is a proton donor while a base is a proton acceptor." Therefore, since aspirin donates a proton (H+), it can be classed as an acid.

Phenolphthalein indicator takes advantage of the fact that aspirin is an acid. In acidic conditions, the phenolphthalein indicator is colourless. However, in basic conditions, the phenolphthalein indicator is pink. The acidic conditions are generated by the aspirin itself resulting in the mixture being colourless when phenolphthalein indicator is added. However, when sodium hydroxide is added, it neutralises the acid by reacting with the H+ to form a neutral solution:

C9H8O­4(aq) + NaOH(aq) è C9H7O4-Na+(aq) + H2O(aq)

As the sodium hydroxide continues to be added, the solution becomes more basic and the phenolphthalein indicator turns pink. This is due to a change in the sub-shell configuration of phenolphthalein resulting in green light being absorbed and the complimentary colour (pink) being emitted [18].

Another way to think of this is that the conjugate acid and base forms of phenolphthalein indicator form different colours.

If we think of phenolphthalein indicator in acidic conditions as HIn we can write equations to demonstrate the different phases that the indicator can be in, as well as the different colours.

An equilibrium reaction may be set up:

HIn(aq) H­+(aq) + In-(aq)

(Colourless) (Pink)

When Sodium hydroxide (NaOH) is added, the OH- group reacts with the H+ ion to produce water:

HIn(aq) + OH-(aq) In­-(aq) + H2O(l)

(Colourless) (Pink)

Since the concentration of H+ has been reduced, the equilibrium shifts towards the products. This causes more In- to be formed and ,as a result, more of a pink colour in the solution[17].

[17] (CI 8.1)

[18] (teachers.sduhsd.net/lcale/Aspirin%20Titration%20Lab.doc)

Thin-Layer Chromatography (t.l.c.) Equipment List

UV light source

Dropping tubes

Silica coated t.l.c. plates

Small beaker

Cling film to cover beaker

Ethanol

Recrystallised aspirin

Pure aspirin sample

2-hydroxybenzoic acid sample

Chromatography solvent - cyclohexane, ethyl ethanoate, ethanoic acid (200:100:1)

Thin-Layer Chromatography (t.l.c) Method

I took a t.l.c. plate and drew a pencil line approximately 1cm from the bottom.

I then took a few crystals of my recrystallised sample of aspirin and dissolved it in a small amount of ethanol. I then repeated this with salicylic acid and pure aspirin.

On the baseline I then placed a small spot of my dissolved aspirin sample along with the dissolved pure aspirin and dissolves salicylic acid spaced out across the baseline. I used a dropping pipette to place the spot.

Working in a fume cupboard, I then poured some of the chromatography solvent into the beaker to a depth of about 5mm. Afterwards, I put my t.l.c. plate into the solvent.

I covered the beaker with cling film and left the solvent to rise up the t.l.c. plate.

When the solvent was close to the top of the plate, I removed the plate and placed it in a fume cupboard to allow the solvent to evaporate.

I then examined the paper under UV light to enhance the spots and make them visible.

[19] (WM 5.1)

Thin-layer chromatography, along with all other types of chromatography, relies on the equilibrium set up when a compound distributes itself between two phases. One phase is called the 'stationary phase' and it does not move. The other phase is called the 'mobile phase.' This mobile phase moves over the stationary phase. In this case, the stationary phase is the t.l.c. paper and the mobile phase is the chromatography solvent.

Different compounds will have different affinities for the mobile or the stationary phase and as a result will move along the mobile phase at different speeds. When the solvent has evaporated off, all the compounds will stop moving along the mobile phase. We can then compare Rf values (see below) to values for known molecules, to determine which compounds we have.

Rf Values

The Rf value for a substance is the distance that substance travels compared to the distance the solvent travels. It is calculated by dividing the distance moved by the solvent by the distance moved by the spot. In the example below, the Rf value for substance A is equal to x/a.

[20] (CI 7.3)

Nuclear Magnetic Resonance (n.m.r.) Spectroscopy

Nuclear Magnetic Resonance (n.m.r.) works differently than mass spectrometry as it uses the behaviour of the nuclei of different atoms rather than the acquired charge of certain molecules such as 1H. These nuclei behave like magnets when they are placed in a magnetic field and some line up with the magnetic field whereas some line up against the magnetic field. If a nucleus lines up against the magnetic field, then it has a higher energy than one that is aligned with the magnetic field. If the correct frequency of radiation is applied, some nuclei will move up to the next energy level and absorb energy. The energy which they absorb corresponds to radio frequency.

The energy that a nucleus needs to absorb to move up to a higher energy level depends on the strength of the magnetic field that it experiences. This magnetic field is affected by other atoms in the molecule as they each have a magnetic field of their own. This means that for every type of molecular arrangement, there is a slightly different magnetic field. The atoms have different energy gaps between their high and low energy levels and so absorb different frequencies of radiation. As a result, they give different n.m.r. absorption peaks and we can find out how many hydrogen atoms of different types there are in a molecule.

[Diagram of n.m.r. apparatus]

[21](http://www.uz.zgora.pl/~jleluk/ppt/lipids2pol/www.cem.msu.edu/_reusch/VirtualText/Spectrpy/nmr/nmr1.htm)

Interpreting an n.m.r. spectrum

Running along the bottom of an n.m.r. spectrum is a tetramethylsilane (TMS) reference which is used as standard reference because it gives a sharp signal well away from most of the substances of interest to chemists. The n.m.r. spectrum is read by the extent at which a signal differs from the TMS reference. For example, the n.m.r. spectrum for ethanal (see below) has a peak at a chemical shift of around 10, and then another peak which is three times higher at a chemical shift of around 2. The peak at 10 is from the CHO part of the molecule whereas the peak at 2 is from the CH3­ part of the molecule. The reason the CH3­ peak is three times larger than the CHO peak is because it has 3 more hydrogen atoms. The integrated trace goes up in steps which are proportional to the areas of absorption signal. This means that we know how many protons are being absorbed each time.

[Diagram of n.m.r. spectrum of Ethanal]

[22](http://www.avogadro.co.uk/analysis/nmr/nmr.htm)

Iron(III) Chloride

Iron(III) (Fe3+) is a transition metal ion and is used in this case to test whether or not my aspirin was hydrolysed back into salicylic acid. A transition metal is defined as:

An element which forms at least one ion with a partially filled sub-shell of d electrons.

In an atom, electrons are arranged in shells which are split into sub-shells: s, p, d, and f. The letters of these sub-shells also correspond to different parts of the periodic table which is why transition metals are also known as 'd-block elements.' Each sub-shell can hold a certain number of electrons:

Sub- shell Maximum number of electrons

s 2

p 6

d 10

f 14

The electronic configuration of elements can also be given in terms of sub-shells. For example for iron (Fe) the electronic configuration can be written as:

1s22s22p63s23p63d64s2.

In this case, each regular number represents a shell, each letter represents a sub-shell and each superscripted number represents how many electrons each sub shell is holding. By summing the values of all the superscript numbers, we can see that an Fe atom has 26 electrons.

The structure of iron can be shown in a table which takes into account only the sub shells. This can then be compared to different ions of iron.

Ion 3d 4s

↑↓

↑

↑

↑

↑

↑↓Fe

↑↓

↑

↑

↑

↑

Fe2+

↑

↑

↑

↑

↑Fe3+

Each arrow represents an electron and the different oxidation states of iron are shown. This allows iron to form a range of different molecules including complex molecules with ligands. Ligand complexes form as dative bonds form with the central metal ion [23].

The Iron(III) Chloride test is used to distinguish between aspirin and salicylic acid. This is due to the fact that salicylic acid has a phenol group whereas aspirin does not. The Iron(III) Chloride reacts with the phenol group to form a ligand complex, and HCl [24]:

FeCl3 + 6C7H6O3 è [Fe(C7H5O3-)6]3-­ + 3H+Cl-

The complex formed has an octahedral shape due to there being 6 areas of electron density around the central Fe3+ ion. This means that there are a mix of 120° and 109° angles. It also has a purple colour which makes it easily distinguishable in solution.

Diagram showing 3D representation of complex:

The lone pair of electrons in the ligand (C7H5O3-) forms dative bonds with the Fe3­+ ions. The lone pair also repels electrons in the 3d orbital of the Fe3+ ions which cause a splitting of the sub-shell. This results in two different energy levels. Therefore, when energy is applied (in this case, light) the electrons are excited and absorb energy in the green-yellow wavelength. When the electrons lose energy again, they emit light with the complementary colour which, in this case, is a purple-violet colour [25].

Diagram showing dative bonds between (C7H5O3-) and Fe3+:

Taking this into account, I put a small sample of my aspirin in a test tube and added a few drops of Iron(III) Chloride. I repeated this test 3 times and recorded my results in the results section.

[Phenol group in Salicylic acid]

[23] (CI 11.5)

[24](http://wiki.answers.com/Q/What_is_the_chemistry_of_the_reaction_between_phenol_and_neutral_ferric_chloride)

[25](http://en.wikibooks.org/wiki/A-level_Chemistry/OCR_(Salters)/Complexes#How_do_ligands_split_d_orbitals.3F)

Risk Assessment: Synthesise Aspirin

Hazard

Probability of occurrence (1=Very Improbable, 10=Very Probable)

Severity of occurrence (1=Non Severe, 10= Very Severe)

Precaution to prevent occurrence of hazard

Ethanoic anhydride is corrosive - could cause blindness or chemical skin burns if spilled

6

7

Wear gloves, goggles and a lab coat while handling.

Take care while pouring.

Concentrated Sulfuric(VI) acid is corrosive - may cause blindness or chemical skin burns if spilled

6

9

Wear gloves, goggles and a lab coat while handling.

Take care while pouring.

Glacial Ethanoic acid is corrosive - may cause blindness or chemical skin burns if spilled

6

7

Wear gloves, goggles and a lab coat while handling.

Take care while pouring.

Oven is hot - burns may occur when retrieving aspirin

4

8

Wear thermal gloves when retrieving substances and keep oven at arms length.

Risk Assessment: Recrystallisation

Hazard

Probability of occurrence (1=Very Improbable, 10=Very Probable)

Severity of occurrence (1=Non Severe, 10= Very Severe)

Precaution to prevent occurrence of hazard

Severe burns from Bunsen burner

5

9

Ensure safety flame is on while lighting Bunsen burner.

Use tools for heating test tubes and use thermal gloves for moving beakers off gauze.

Boiling water may cause skin burns

6

8

Wear goggles and a lab coat while heating water in Bunsen burner.

Hot solution spitting out of beaker over Bunsen flame - may cause severe burns and blindness if it spits in eye

7

9

Wear lab coats and thermal gloves. Ensure you wear goggles when near solutions over Bunsen flames.

Risk Assessment: Aspirin Assay

Hazard

Probability of occurrence (1=Very Improbable, 10=Very Probable)

Severity of occurrence (1=Non Severe, 10= Very Severe)

Precaution to prevent occurrence of hazard

Ethanol is flammable which may result in serious burns. It is also irritating to skin and eyes.

4

7

Do not handle ethanol near flames.

Always wear goggles, gloves and a lab coat.

Phenolphthalein is an irritant and may cause skin rash or burning

6

7

Wear gloves, goggles and a lab coat while handling phenolphthalein.

Sodium hydroxide is corrosive and may cause skin burns.

6

8

Wear lab coats, gloves and goggles when handling Sodium hydroxide.

Risk Assessment: T.L.C.

Hazard

Probability of occurrence (1=Very Improbable, 10=Very Probable)

Severity of occurrence (1=Non Severe, 10= Very Severe)

Precaution to prevent occurrence of hazard

Chromatogaphy solvent is highly flammable and an irritant to the respiratory tract

3

9

Do not handle solvent near flames.

Always wear goggles and a lab coat.

Glassware (beaker, capillary tubes) may cause harm if broken.

5

8

Wear goggles and a lab coat while handling glassware.

Prolonged exposure to UV may cause blindness, skin burns and cancer.

6

9

Face UV light away from you and others at all times. Ensure goggles are on when using and do not look directly at the light.

Risk Assessment: Melting point

Hazard

Probability of occurrence (1=Very Improbable, 10=Very Probable)

Severity of occurrence (1=Non Severe, 10= Very Severe)

Precaution to prevent occurrence of hazard

Severe burns from Bunsen burner

5

9

Ensure safety flame is on while lighting Bunsen burner.

Use tools for heating test tubes and use thermal gloves for moving beakers off gauze.

Glassware (capillary tubes) may cause harm if broken or heated

5

8

Wear goggles and a lab coat while handling glassware.

Place glassware on a heatproof mat to cool after melting.

Hot apparatus may cause burns

5

9

Allow time for apparatus to cool before touching it.

Results: Iron(III) Chloride Test

Compound

Observation after adding Iron(III) Chloride

Recrystallised Aspirin

No Change

Pure Aspirin

No Change

Salicylic Acid

Turns purple

Results: Aspirin Assay

My Recrystallised Sample

Rough

1st

2nd

3rd

Initial Reading* (cm3)

0.00

7.55

4.50

5.65

Final Reading* (cm3)

16.30

23.90

20.75

21.95

Titre (cm3)

16.30

16.35

16.25

16.30

Average Titre

16.30

Pure Aspirin

Rough

1st

2nd

3rd

Initial Reading* (cm3)

0.00

2.35

1.55

5.80

Final Reading* (cm3)

16.35

18.75

17.90

22.10

Titre (cm3)

16.35

16.40

16.35

16.30

Average Titre

16.35

Aspirin Tablet

Rough

1st

2nd

3rd

Initial Reading* (cm3)

0.00

3.60

2.65

5.35

Final Reading* (cm3)

14.40

17.95

17.10

19.75

Titre (cm3)

14.40

14.35

14.45

14.40

Average Titre

14.40

*of Sodium Hydroxide volume

Results: Thin-Layer Chromatography

Substance

Distance moved (mm)

Rf value

Solvent

65.0

N/A

Pure Aspirin

27.5

0.42

Recrystallised Aspirin

28.0

0.43

Salicylic Acid

43.0

0.66

Results: Melting Point

Substance

Start Melting (°C)

End Melting (°C)

Melting Range (°C)

Salicylic Acid

158.5

159.0

0.5

Pure Aspirin

135.0

135.0

0.0

Recrystallised Aspirin

135.0

136.0

1.0

Analysis of Results: Aspirin Assay

For each of my samples, I will use the average titre to determine the purity. This is because the average titre is a useful indicator of the overall purity of the entire sample. I will use the equation:

C9H8O­4(aq) + NaOH(aq) è C9H7O4-Na+(aq) + H2O(aq)

And the mole equations:

n=c*v, n=m/RMM

I will divide each sample of aspirin which I measured by 5 as I made up a standard solution which would allow me to do 5 titrations. However, after gaining four concordant results I concluded that it was safe to assume that the amount of sodium hydroxide used would be sufficient to provide an accurate assay.

My aspirin

Pure Aspirin

Aspirin tablet

1. Mass of aspirin powder used - g

1.527

1.477

1.462

2. (Mass of aspirin powder used)/5 - g

0.305

0.295

0.292

3. Volume of sodium hydroxide used - cm3

16.30

16.35

14.40

4. Amount of sodium hydroxide used - mol (c*v)

1.63x10-3

1.635x10-3

1.44x10-3

5. Amount of aspirin which reacted with sodium hydroxide - mol

1.63x10-3

1.635x10-3

1.44x10-3

6. Mass of aspirin which reacted with sodium hydroxide - g (n*RMM)

0.293

0.294

0.259

7. Aspirin in powder used - % (Step6*100/Step2)

96.1%

99.7%

88.7%

[16] (WM 6)

The solution turns pink due to changes in energy levels and indicator theory as mentioned previously (pgs 14-15). When the HIn dissociates to form H+ + In-, and OH- groups react with the H+ group to form water, the equilibrium is shifted to the products and more of the pink In- is formed.

From my results, we can see that the value for my recrystallised aspirin is 3.4% less pure than the value for the pure aspirin. This could be due to impurities which may not have been extracted during the recrystallisation. As a result there would be less aspirin and, therefore, less acid to react with the NaOH. This means that the end point of the reaction will be reached with less NaOH being used.

The aspirin tablet, on the other hand, has other additives which mean that the purity is lower. This results in even less acid available to react with the NaOH. Therefore, the end point is reached with a lower amount of NaOH which results in the percentage purity being low even though the aspirin used in the tabled may be pure.

Ananlysis of Results: Thin-Layer Chromatography

Since the Rf values for thin-layer chromatography differ with different solvents, there is no published data with set values. This means that I had to compare my results to pure aspirin and salicylic acid. From my results, I found that the Rf value of my recrystallised aspirin was very similar to the value for the pure aspirin. It was also very far from the value for salicylic acid. This means that my sample of aspirin was very close in purity to pure aspirin.

The salicylic acid moved further up the chromatography plate than the aspirin as it has a lower affinity for the plate and a higher affinity for the solvent. This means that it is more likely to travel with the solvent than to settle on the plate (see pg 17).

Each sample has a different affinity for the mobile stage (see pg 17), and so each sample will move different distances from the base line and therefore have a different Rf value. If the Rf values for my aspirin sample and pure aspirin are similar then we can assume that the level of purity is relatively similar.

Analysis of Results: Iron(III) Chloride

Comparison of the results of pure aspirin and salicylic acid shows that pure aspirin showed no change when Iron(III) chloride was added, however, salicylic acid turned purple. When I added Iron(III) chloride to my sample of recrystallised aspirin, I found that it showed no change. Therefore I could assume that did not have phenol impurities.

The phenol group in the salicylic acid will go purple if Iron(III) chloride is added because a complex is formed with oxidised salicylic acid and Fe3+. This causes a splitting of the d sub shells in the ligand and as a result excited electrons absorb the green wavelength. This causes the complimentary colour (purple) to be emitted.

However, this does not occur with aspirin so if we have aspirin then the colour will remain orange.

Analysis of Results: Mass Spectrometry

Results have not been returned.

Analysis of Results: Nuclear Magnetic Resonance Spectroscopy

Results have not been returned.

Analysis of Results: Melting Point

After analysing my results, I found that the point at which my sample of aspirin melted was very close to the point at which pure aspirin melted. The pure aspirin melted at the published melting point for aspirin which shows that my results would have been accurate. However, the melting range for my recrystallised aspirin was larger than the melting range for the pure aspirin. This means that it is less pure than the pure aspirin. This is because the melting points of impurities such as salicylic acid are higher than the melting point of aspirin itself. This means that the point at which the aspirin melts will be lower than the point at which any salicylic acid melts which would result in a larger melting range.

Different intermolecular bonds in aspirin and salicylic acid cause them to melt at different temperatures. Aspirin has one area for hydrogen bonding and is a relatively large molecule whereas salicylic acid has two areas for hydrogen bonding and is relatively similar in size to aspirin. Both molecules have areas of permanent dipole-permanent dipole bonds. Since both molecules have permanent dipole - permanent dipole bonds but salicylic acid has two areas of hydrogen bonding rather than one, the melting point of salicylic acid will be higher than that of aspirin because hydrogen bonds are the strongest form of intermolecular bonding. It also means that the melting points of substances are very specific and, therefore, if my sample of aspirin melts at the published melting point then we can assume that it has a high level of purity.

Evaluation and Limitations : Synthesise Aspirin

Percentage Uncertainty

Percentage error is calculated as the (degree of precision/value)*100. Using manual equipment e.g. a measuring cylinder, measurements can be precise to a degree of 1/2 of the smallest increment. For example the measuring cylinder measures in ml; therefore the value is precise to a degree of +/-0.5 ml. The percentage uncertainty, therefore, of the measurements I took from the measuring cylinder are:

(0.5/20)*100 = 2.5%

This applies to both measurements as I used 20cm3 of Ethanoic anhydride and glacial Ethanoic acid. This means that overall; the measurements I used for my solutions had a total of a 5% possibility of error.

However, for digital equipment such as the scales which I measured the Salicylic acid and Aspirin on, the measurements can be precise to a degree of +/-½ of the smallest unit. For example, the scales I used were accurate to 3 decimal places which means that the values I took are precise to +/-0.0005g. Therefore, the percentage uncertainty of the measurements I took from the scales is:

(0.0005/10.053)*100 = 0.005% and (0.0005/11.006)*100 = 0.005%

This means that the error of measurements for my Salicylic acid and final Aspirin sample is 0.01%.

Measuring cylinders require the student to make a judgement of when the bottom of the meniscus is at the marking on the cylinder. This leads to great inaccuracies as it is not digital and there is a lot of room for human error. The measurements from the measuring cylinder are precise to a degree of +/-0.5 ml. The scales, on the other hand, were digital and measured to 3 decimal places. This means that the values of solid substances can be accurately measured with a precision of +/-0.0005g. Also, while drying the sample of aspirin, some liquid solution may still be trapped. This would change the mass of the final measurement of the aspirin sample.

[26]www.ece.rochester.edu/courses/ECE111/error_uncertainty.pdf)

Evaluation and Limitations: Purify Aspirin

While purifying my sample of Aspirin, I had to ensure that I did not heat for too long and that the temperature didn't get too high. This is because heating for a long time at high temperatures would cause the Aspirin to be hydrolysed back into Salicylic acid due to the fact that a 10°C increase in temperature causes the rate of reaction to double which increases the chance of hydrolysis of salicylic acid. However, it is the method used for purifying organic solids. The fact that soluble and insoluble impurities are filtered out make it very effective.

Evaluation and Limitations: Iron(III) Chloride

The Iron(III) Chloride can be used to tell if I have Aspirin or not but it wouldn't tell me how pure it is. However, this test is not specific to Salicylic acid and Aspirin so it is not 100% reliable as I could have a completely different molecule to aspirin and it still wouldn't turn purple due to the fact that a new compound made might not have a phenol group.

Evaluation and Limitations: Aspirin Assay

From my results, I found that the value of purity of aspirin from the supermarket was 88.7%. Also, after measuring the tablet at 337mg, and reading on the packaging that each tabled contained 300mg of aspirin, I found that the aspirin content in each tabled is actually 89.0%. Assuming that I was consistent with all of my titres, this means that my results were reasonably accurate. However, the Sodium hydroxide I used may not have been exactly 0.1 mol dm-3. If the concentration of Sodium hydroxide was any lower than 0.1 mol dm-3 then the values of the titres that I measured would be higher than the values of the titres for if it was exactly 0.1 mol dm-3. This would impact my calculations and result in the conclusion that the aspirin samples were more pure than they actually are. Another action which may have impacted my results is the idea that I could have judged the end point too late. If this was the case, I would have a higher value for the titre and found that the aspirin was more pure than it actually was.

Percentage uncertainty could also have a role. Since I was using a burette, I could only be precise to a degree of +/-0.05 ml. This value would have to be multiplied by two as I am taking two readings. I will use the average titre to calculate the percentage error from [26]. This means that I use the formula:

2(0.05/titre)*100

I also had to calculate the error for the weight of my recrystallised aspirin, the pure aspirin and the salicylic acid in each titre. This is calculated by using the formula:

(0.0005/weight)*100

The weight used in each titre is my original weight divided by 5 as a standard solution was made up which allows 5 titres for each sample. I then added the two errors to get a total error for each titre.

Using these values I calculated the percentage error for each of my average titres:

Sample

Percentage Error

Recrystallised Aspirin

0.77

Pure Aspirin

0.78

Aspirin Tablet

0.86

The percentage errors from my titrations did not exceed 0.86%. This is a very low error and would not greatly impact my results. This means that this is an accurate way to measure the percentage purity of my aspirin samples.

Evaluation and Limitations: Thin-Layer Chromatography

Thin-layer chromatography gives us an idea of the purity of my aspirin compared to salicylic acid and pure aspirin. However, it does not give percentage purity. Also, since the chromatography paper is small, a small change in Rf ­value could mean a large change in purity. Human error could also be made when measuring the distance moved by each spot which could result in a different Rf value.

Due to the fact that I used a ruler to measure the Rf value, my measurements could only be precise to a degree of +/-0.5mm. This means that the purity I found may not correspond to the true value of the purity of my aspirin [26]. Using the formula:

Percentage Error = (0.5/distance)*100

I calculated the error for the values that I obtained for the distance moved by each spot and solvent in the t.l.c.:

Substance

Error (%)

Pure Aspirin

1.8

Recrystallised Aspirin

1.8

Salicylic Acid

1.2

The largest error came from my sample of pure aspirin at 1.8%. This is not a large error which means my values Rf values are relatively precise. In conclusion, this is an accurate way to qualitatively compare the purity of my aspirin to that of pure aspirin and salicylic acid. Salicylic acid is used as it is the most likely impurity to arise.

Evaluation and Limitations: Melting Point

Similar to t.l.c., melting point gives a comparison of purity with salicylic acid and pure aspirin, but it does not give a percentage. It also relies on human judgement for the start, and end, of the melting process. This could lead to errors of measurements as it is not done digitally.

As the thermometer is measured and read manually, I can only be precise to a degree of +/-0.5°C. This means that the melting range may not be as precise as I measured meaning the purity of my aspirin may not correspond to the true value of purity for my aspirin [26]. To measure the percentage error, I have to use:

(0.5/melting point)*100

I have to do this for the original and final melting values. The results which I obtained for percentage error are shown below:

Substance

Start Melting

End Melting

Total

Salicylic Acid Error (%)

0.32

0.32

0.64

Pure Aspirin Error (%)

0.37

0.37

0.74

Recrystallised Aspirin Error (%)

0.37

0.37

0.74

The errors for my results do not exceed 0.74% which is not a large error. This means that, although my results may be slightly imprecise, they will not be impacted greatly. Therefore, I can conclude that this is an accurate way to qualitatively compare the purity of my aspirin to that of pure aspirin and salicylic acid. Salicylic acid is used as it is the most likely impurity to arise.