A Level Chemistry Project- the Purity of Aspirin

To research, using various sources, the history of aspirin, its use in medicine, methods of synthesizing it and of measuring its purity – To compare the % purity of a branded aspirin tablet with a generic aspirin tablet – To compare 2 methods of composition analysis of the two types of aspirin Research The history of Aspirin 1 400BC, Greece- Hippocrates recommended the use of a brew made with willow leaves to help with the pain of childbirth. 763, England- Edward Stone, a clergyman, having read a paper regarding the use of Willow bark on malaria (Agues), collected observations from around the country on its effects in relieving the fever that comes with malaria. 1830’s- a Scottish physician found that willow bark relieved symptoms of acute rheumatism. 1840’s- Organic Chemists identified the active ingredient in willow bark to be salicin. 1874, Dresden, Germany- Salicylic acid was first made and sold as a pain-killer, but it caused severe irritation of the mouth, gullet and the lining of the stomach. Salicylic acid (2-Hydroxybenzoic acid) 875- Chemists made sodium salicylate for doctors to try on patients. It was found to have fewer digestive side-effects, and still worked to reduce pain, however it tasted horrible and caused rheumatism patients to vomit. 1890’s, Germany- Felix Hoffman of Bayer pharmaceuticals made aspirin with good medicinal properties, low membrane irritation and a reasonable taste. The name ‘aspirin’- a-acetyl, as the systematic name for aspirin was acetylsalicylic acid spir- spirea plant which yields salicin Asprin (2-Ethanoyloxybenzenecarboxylic acid) 1898- Aspirin was sent for clinical trials, Bayer manufactured the medicine and patented the process. 915- The British needed aspirin in WW1, but since it was made by the Germans, the government offered a ? 20,000 reward to whoever could develop a workable manufacturing process. George Nicholas, a Melbourne pharmacist, did this and called his tablet ‘Aspo’. 1982- John Vare, who was that year receiving a Nobel prize for his work on prostaglandins, discovered aspirin and some other pain-killers/anti-inflammatory drugs to inhibit the key enzyme in the prostaglandin pathway, thereby stopping the production of some prostaglandins that cause pain and inflammation.

Now types of aspirin are made all over the world each year, and it is not only used as a painkiller but is also effective in reducing the incidence of heart disease. Chemistry of aspirin and its use in medicine 2 * Chemical structure of acetylsalicylic acid Name | acetylsalicylic acid (aspirin) | Systematic name | 2-acetoxybenzoic acid | Formula | C6H4(OCOCH3)CO2H | Melting point | 136°C (277°F) | Boiling point | 140°C (284°F) | Molecular mass| 180| * * Aspirin is an aromatic acetate aromatic: organic chemical made by natural sources containing a benzene ring -its structure also includes an ester group and a carboxylic group attached to adjacent carbon atoms on the benzene ring * Aspirin is a derivative of salicylic acid -salicylic acid: C7H6O3, Mr: 138 -contains a phenol group, the hydroxyl of which is situated in the same place as the ester group in aspirin. -there’s also an adjacent carboxyl group * Preparation of salicylic acid: In Kolbe synthesis, a common industrial method of manufacturing salicylic acid for aspirin, sodium phenoxide is heated under pressure with carbon dioxide.

The reaction mixture is then acidified to produce salicylic acid. 3 * Preparation of aspirin: esterification of the phenolic hydroxyl group of salicylic acid 3 Once synthesised, if aspirin is allowed to come into contact with moisture, hydrolysis occurs, as with all esters, producing ethanoic and salicylic acid. This enables autocatalysis to occur; the hydrolysis of esters is catalysed by acids, and both products are acids. Hence hydrolysis is slow to start with before accelerating as more hydrogen ions are produced Use in medicine: action mechanism 2

Aspirin has analgesic, antipyretic and anti-inflammatory properties. Aspirin works by limiting the body’s production of prostaglandins, which are fatty acid derivatives causing inflammation and pain. They do this by irreversibly inhibiting cyclooxygenase enzymes which produce prostaglandins and thromboxane . This is why aspirin is commonly used for relief from headaches, muscular pains, arthritis and menstrual pain. It is also used to reduce the risk of heart attacks and strokes due to thromboxane stimulating the clotting of platelets. A prostaglandin A thromboxane

In commercial aspirin products, a small amount (about 300mg) of acetylsalicylic acid is bound together with a starch binder and sometimes buffers and caffeine to form a tablet. In the small intestine, the acetylsalicylic acid is broken down to give salicylic acid, which is absorbed by the bloodstream. The function of the buffer is to reduce any irritation that might be cause by the carboxylic acid group of the aspirin. Choosing methods of composition analysis Thin layer chromatography is one method of separating and identifying aspirin.

The procedure is carried out by placing small amounts of synthesised product, salicylic acid and commercialised aspirin, on one edge of a chromatography plate. When placed in a container with the solvent, the solvent will flow up the chromamtogram carrying the samples with it. The differing solubilities of each sample in the solvent means that they will travel a different distance up the plate (more soluble will move higher, more attracted to the plate will remain nearer to the line). UV light will then be used to detect the aspirin after removing the plate.

Small amounts of inert binder will show manufactured aspirin tablets to be less than 100% purity. 4 Although this method can be used to identify the presence of aspirin, it does not allow quantitative analysis of the purity if aspirin, which is why I won’t pursue this method. Another unsuitable option of analysis is by melting points. I will not use this method because it requires the removal of all starch, silica and cellulose from the tablets. Many processes such as suction filtration and crystallisation would be required just to isolate the aspirin. It would therefore not be a worthwhile method in the time allowed.

Other disadvantages are that some of the starch may still accompany the aspirin into the filtrate, and the fact that it’s unlikely that the active ingredient, salicylic acid, will actually crystallise out at the same time as the aspirin, due to its higher solubility. Both of these factors mean that the method is unlikely to be successful. A more suitable method is back titration. This is likely to be reliable due to taking into account the fact that an impurity in aspirin is salicylic acid, making it accurate. In this way it is a better method to standard titration (see precursors of back titration)

Colorimetry would be a good method for comparison with back titration, as the two are very different. The back titration requires alkaline hydrolysis of the aspirin, whereas colorimetry requires several separate samples undergoing acid hydrolysis. Colorimetry will directly measure the concentration of salicylic acid in a hydrolysed aspirin solution, whereas back titration calculates the initial amount of aspirin present via a method of differences. Another advantage to the colorimetry is that it requires little apparatus, so is a suitable method under the given conditions and time restraints.

Precursors: method 1: Back titration Alkaline hydrolysis 7 Aspirin undergoes very weak hydrolysis, due to one aspirin molecule reacting with 2 OH ions. To overcome this, a known excess of base is added to the sample solution. A titration is then carried out with hydrochloric acid to find out the amount of unreacted base. A simple subtraction will then indicate the amount of base which has reacted with the aspirin, and thus the amount of aspirin in the sample. Hydrolysis is a process in which a compound is subject to nucleophilic attack by either an OH ion or water, causing decomposition.

The conditions for this reaction usually include an acid catalyst. As an acid and an ester, aspirin can be hydrolysed to form salicylic acid (2 hydroxybenzoic acid) and ethanoic acid. In the alkaline hydrolysis of aspirin it is the OH ion which performs the nucleophilic attack on the carbonyl carbon to make an intermediate which subsequently loses an acetate ion (nucleophilic acyl substitution). CH3COO. C6H4. COOH + 3NaOH ? NaO. C6H4. COONa + CH3COONa + 2H20 Hydrolysis step 1: HYDROLYSIS OF THE ESTER LINK CH3COO. C6H4COOH + 2NaOH?? NaO. C6H4. COONa +H20 Hydrolysis step 2: REACTION OF PHENOLIC -OH

Having hydrolysed the aspirin, the back titration is carried out using dilute hydrochloric acid. NaOH + HCl ? NaCl + H2O – The excess alkali is destroyed NaOC6H4 +HCl ? HOC6H4COONa + NaCl – Phenoxide ion gains a proton and changes into a phenol group A phenolphthalein indicator will be used which will stop the process at the step. Otherwise, with greater quantities of acid and at lower PH’s, the salt is converted into its parent acid. Preparation and standardization of solutions for the back titration * First the mean mass of both a generic and the manufactured aspirin tablets must be found.

This is done most accurately by measuring the mass of (minimum) 5 tablets at once, and dividing by 5. By not measuring each one individually, measurement errors are proportionally less due to the greater mass. | Mass of 5 tablets g| Average mass of 1 tablet g| Branded aspirin tablets| 2. 155| 0. 4310| Generic aspirin tablets| 3. 008| 0. 6016| * To make up a solution of aspirin of suitable concentration, its solubility must be considered. Given its low solubility, I was advised by a teacher that 0. 05moldm-3 would be an appropriate concentration.

Attempting a higher concentration could be unsuccessful as not all of the aspirin would dissolve. 10 1-Making up the aspirin solution Calculating the mass of aspirin required for a 0. 05moldm-3 solution For all measurements of mass taken, recordings are to 3 decimal places (degree of accuracy) Mr of aspirin = 180. 160 (3dp) Volume required = 500 cm3 Mass of aspirin = 180. 160 ? 500 20 1000 = 4. 504g * However, since aspirin tablets are not 100% aspirin, the amount of tablet required to provide this much must be calculated. Calculating the % of aspirin in a tablet

Xg represents the mean mass of an aspirin tablet Yg represents the mass of the aspirin in that tablet % Aspirin in one tablet = Y ? 100 X | Generic aspirin tablet| Branded aspirin tablet| X g| 0. 6016| 0. 4310| Y g| 0. 300| 0. 325| % aspirin in 1 tablet| 49. 867 | 75. 406| * Therefore the mass of tablet needed to give 4. 505g of aspirin… Mass of tablet needed = mass of aspirin needed X 100 % aspirin in 1 tablet Mass of Generic aspirin tablet needed(g)| Mass of Branded aspirin tablet needed(g)| 9. 034 | 5. 973| Calculating the number of tablets required

Having calculated the mass of aspirin required and thus the mass of tablet required to make up the appropriate solution, for both branded and generic tablets, it must then be calculated how many of each tablet must be used. No. of tablet required = mass of tablet needed Mean mass of 1 tablet Number of generic tablets required| Number of branded tablets required| 15. 016 so 16 | 13. 858 so 14| * The next step is to grind up the given number of tablets using a pestle and mortar. For each type of aspirin the tablets should be ground to the same consistency, and to increase the rate of dissolving in ethanol this hould be as fine as possible. 10 * A container should then be placed onto an electronic balance and the value tared, before the required mass of ground aspirin tablet is measured out onto the scale using a spatula. This is then transferred to a 250ml beaker (ethanol rinsed to remove risk of contamination with impurities) before measuring the mass of the original container again. This is to account for any inaccuracy due to residual powder remaining in the container: Mass transferred = original mass (before transfer) – final mass(after transfer) | Generic aspirin| Branded aspirin|

Original mass (g) (initial tared mas)| 9. 041| 5. 974| Final mass (g) (tared mass of receptacle with residual powder)| 0. 094| 0. 125| Mass transferred (g)| 8. 947| 5. 849| Mass of aspirin = mass of powdered tablet transferred X proportion of aspirin in tablet | Generic| Branded| Mass of aspirin (g)| 4. 462| 4. 410| * Ethanol is then added to the beaker to dissolve the aspirin. About 150 cm3 should be required but more may be added if necessary, bearing in mind some possible components of an aspirin tablet eg lactose, are highly insoluble.

If a vortex forms from undissolved powder, it should be gently tapped with a glass rod to aid the dissolving process. * I will then use 2 500cm3 volumetric flasks to contain the generic aspirin solution and the branded aspirin solution. The transfer will be done using a funnel and pouring each solution down a glass rod into the flask, ensuring no solution and thus aspirin is lost due to splashing * The beaker and glass rod should then be washed with ethanol several times. After each rinse the ethanol is then emptied into the volumetric, once again to ensure that no residual aspirin is lost.

The funnel should also be rinsed with ethanol before it is removed (when the solution reaches the neck of the flask), and ethanol added slowly with a teat pipette when it nears the graduation mark. * The neck of the flask should then be wiped with paper to ensure no change occurs to the volume of the solution caused by drips on the edge of the flask. * Each flask will then be stoppered securely and inverted a few times to ensure even distribution of molecules in solution. Calculating the actual concentration of aspirin solution Concentration of aspirin = mass of aspirin g ? olume Mr = mass of aspirin(g) ? 0. 5 dm3 180. 16 Concentration of generic solution (moldm-3 )| Concentration of Branded solution (moldm-3 )| 0. 0495| 0. 0490| 2i)- Making up the sodium hydroxide solution * For this experiment the required concentration of NaOH is 0. 1moldm-3 . The supplied solution has concentration of 2M, so must be diluted. I require 2dm3 volume of NaOH solution. Dilution factor = concentration of current NaOH solution Concentration of required NaOH solution = 2. 0 0. 1 = 20 Calculating the amount of NaOH solution needed to make 2dm3 = volume of 0. moldm-3 NaOH solution needed Dilution factor = 2. 0 20 = 0. 1dm3 This amount of solution is accurately measured out using a pipette * Subsequently a 2dm3 volumetric will be rinsed with distilled water to reduce risk of contamination with impurities. It is then partially filled with distilled water before the pipette with NaOH solution is emptied into the volumetric. This is done carefully by ensuring that the pipette touches the surface of the water before emptying it, to ensure its full volume is transferred * Then distilled water fills the rest of the volumetric, making up 2dm3 .

The inner-neck is dried and stoppered securely, and the volumetric inverted, as in the aspirin solution. 2ii) Standardising sodium hydroxide This step is actually preceded be steps 3i) and 3ii) as a standardised solution of hydrochloric acid is required to standardise the sodium hydroxide solution * Standardised by being titrated against some standardised HCl produced in step 3ii) * The burette is rinsed with some of the hydrochloric acid of known concentration, before filling it with a suitable volume of HCl which does not exceed the highest marker. 20cm3 of the unknown, NaOH, is then pipetted into a distilled water rinsed conical flask. A few drops of phenolphthalein indicator, since a strong acid is being titrated against a strong alkali, are added to the flask. This indicator is slightly acidic, hence only a few drops are added so as not to effect the pH by adding extra H+ ions * Titrate, recording the volume of HCL before and after. The end-point colour change will be from pink to colourless. Having performed 1 rough titration, do at least 3 more to obtain 3 volumes within 0. 1cm3.

Each time when nearing the end-point, add the acid drop by drop. * Take the mean of the 3 volumes within 0. 1, and use this to calculate the concentration of NaOH: 19. 4 cm3 + 19. 3 cm3 +19. 4 cm3 = 19. 367 cm3 of standardised HCL required 3 Moles of HCl = Volume used in titration (cm3 ) X concentration (found in step 3) =0. 019367dm3 X 0. 098 =0. 001897966 * Having found the number of moles of HCl which have reacted we know the number of moles of NaOH given they react in a 1:1 ratio Calculating the actual concentration of the NaOH solution

Concentration NaOH= number of moles volume = 0. 001897966 0. 02dm3 =0. 095 moldm-3 (3dp) 3i)- Making up the hydrochloric acid solution(1) * Once again a solution of concentration of 0. 1moldm-3is required. Since the supplied concentration of HCL is 1M, dilution is required again. Dilution factor = concentration of current HCl solution Concentration of required HCl solution = 1. 0 0. 1 = 10 Calculating the amount of HCl solution needed to make 1 dm3 = volume of 0. 1moldm-3HCl solution needed Dilution factor = 1. 0 10 = 0. dm3 * As with the NaOH solution, a volumetric is then rinsed for impurities with distilled water, before being partially filled with distilled water (to ensure the pipette can be used to transfer the HCl in an accurate manner) * Then 0. 1 dm3 of 1M HCl is transferred accurately with the pipette, once more ensuring it touches the surface of the water before the solution is released (so the full volume is definitely transferred) * The rest of the volume of the volumetric flask is made up with distilled to 1dm3, and the flask will once again be dried, stoppered and inverted. ii) Standardising HCl (1) * Standardising allows you to ensure that a concentration of a solution is know. In this experiment, I must know the exact concentration of the hydrochloric acid so that the number of moles of sodium hydroxide used in the titration can be calculated. * Standardisation of HCl Is done by titrating it with standard sodium carbonate METHOD 1. Set up retort stand with a burette holder 2. Rinse the burette with a small amount of the HCl which is to be standardised- preventing the acid being contaminated by residue in the burette, changing its concentration 3.

Ensuring that the tap is closed, transfer some HCl into the clamped burette using a funnel. Remove air bubbles by briefly opening the tap, also ensuring the volume is lower than the highest marker. 4. Pipette 20cm3 of sodium carbonate solution into a 250cm3 conical flask (which has previously been rinsed with distilled water to remove contaminants) 5. Add methyl orange to the conical flask (3-4 drops, any more may affect pH as methyl orange is acidic). Methyl orange is suitable because a strong acid and a weak alkali are being used, and in these conditions methyl orange is best indicator to show the most accurate end-point. . Record the exact volume of HCl in the burette before beginning the titration. Control the release of acid with one hand whilst swirling the sodium hydroxide solution with the other, ensuring thorough mixing of acid and base. 7. At the end point, when all of the base has been neutralised, the solution in the conical flask with turn from orange to salmon-pink. 8. When nearing this end point, acid should be added drop by drop. 9. Record the volume of HCl in the burette again, and work out the amount used 10. Having done 1 rough titration, repeat the process 3 times more to obtain results within 0. 1cm3

Calculating the exact concentration of HCl -Take a mean of the three volume values which were within 0. 1 cm3 of each other for HCl = (41. 2cm3 + 41. 2cm3 + 41. 1cm3 ) / 3 =41. 167 cm3 (3dp) Concentration of HCl = Number of moles of HCl Volume of HCl used -Before this equation can be used the number of moles of HCl must be calculated moles of Na2CO3 = concentration of Na2CO3 X volume of Na2CO3 =0. 1008409 X 0. 02dm3 =0. 002016818 moles since HCl and Na2CO3 react in a 2:1 ratio: moles of HCL = 0. 004033636 moles Concentration of HCL= moles ? volume =0. 004033636 0. 041167 =0. 098 moldm-3 (3dp) – Making up sodium carbonate solution of known concentration * Required for the standardisation of HCl * To do this, its concentration must be known, so I will make up a solution of concentration 0. 1moldm-3 * So I will need to calculate the mass of Na2CO3 required to make 500 cm3 of this concentration Molecular mass of Na2CO3 = 105. 98874 g * To make up 500cm3 of a 0. 1 moldm-3 solution this value must therefore be divided by 20 to get 5. 299 (to 3 dp) * A container will be placed on the balance, and the value will be zeroed. I will add sodium carbonate to as close to 5. 99 as I can get. This is then transferred to a 500cm3 beaker which has be de-contaminated by rinsing it with distilled water. The container is then weighed again to account for residual sodium carbonate in it. The weight can be subtracted from the original weight to find the exact amount in the beaker Initial tared mass of sodium carbonate – tared mass of receptacle and residual powder = mass of sodium carbonate used 5. 372 – 0. 028 = 5. 344g * Distilled water will then be added to the beaker (150ml aprox) and a magnetic stirrer is used to help dissolve the sodium carbonate.

Any vortex formed will once again be tapped with a glass rod to encourage dissolving. * The solution should then be funnelled into a 500cm3 volumetric flask, by pouring it down a glass rod to avoid splashing and so loss of sodium carbonate. * The beaker and glass rod should then be rinsed several times with distilled water, this then poured into the volumetric flask so that all dissolved solid is definitely transferred. The funnel should also be rinsed with distilled water before it is removed from the beaker, when the solution reaches the neck of the flask. Then more distilled water is added using a teat pipette to make the solution up to 500cm3 exactly, where the bottom of the meniscus touches the graduation mark. To keep this constant, the inside of the neck of the flask should be wiped with absorbent paper to catch any drips which could increase the volume. * The volumetric flask should be stoppered and inverted several times to ensure that the molecules are evenly distributed. This is the standard solution. Calculating the actual concentration of the sodium carbonate solution produced Concentration = no. of molesNo. of moles= mass olume Mr =5. 344g 105. 98874 =0. 0504 moles =0. 0504 0. 5dm3 =0. 1008 moldm-3 Precursors: Method 2: Colorimetry This time, acid hydrolysis of the aspirin tablet is required. This involves the nucleophilic attack of water with an acid catalyst: CH3COOG + H20 ?? H0. C6H4COOH + CH3COOH? It is the property of salicylic acid to from a purple complex with iron (III) ions which makes colorimetry a suitable method of quantitative composition analysis of aspirin; a yellow-green filter can be used to measure the intensity of the purple.

The concentration of salicylic acid in hydrolysed aspirin solution can be found, and hence the concentration of aspirin in the original solution, from a calibration graph made by comparing a range of salicylic acid concentrations combined with a source of iron (III) ions such as iron (III) chloride. 9 The use of iron (III) chloride explains why I will use acid rather than alkaline hydrolysis; it must be acidified to form a complex with the salicylic acid. OH ions wouldn’t work as they would form a precipitate with iron (III) chloride.

The beer-lambert law is what is used in colorimetry to calculate the concentration of a solution. How colorimetry works is that light of a specific wavelength is shone through the sample and detected on the other side. The amount of light absorbed or transmitted can be recorded. In this case the values for absorbance are required for the calculation. 9 The beer-lambert law: A = ?? c A= absorbance ?= absorption coefficient ?= path length 8 c= concentration The calibration graph created will show A plotted against c for known concentrations of salicylic acid with excess iron (III) chloride.

According to the formula it must show a straight line though the origin. Preparation of chemicals for colorimetry 1. Making up salicylic acid * Salicylic acid is required to form a complex with iron (III) ions and subsequently used to create a calibration graph for the colorimeter. A wide range of concentrations must be created so as to encompass the likely concentration of salicylic acid in the dissolved aspirin samples. * For hydrolysis I made up aspirin solutions of 0. 05 moldm-3, which were then diluted. Calibration solutions must therefore be produced ranging from 0. 00 to 0. 01- moldm-3 * First I need a starter solution of salicylic acid of concentration more the 0. 01 moldm-3, so that it can be diluted, and of suitable quantities that It can be used for each concentration of solution and not run out (since 10 concentrations will be required to make a suitable calibration graph) – I plan to make up 1 dm3 of 0. 04 moldm-3 concentration salicylic acid solution Calculating the amount of salicylic acid powder required Number of moles = volume X concentration =1 dm3 X 0. 04 moldm-3 =0. 04 moles

Mass (g) = Number of moles X Mr (of salicylic acid) =0. 04 X 138. 123 =5. 525g required – This mass of salicylic acid powder is measured out as previously described, doing the initial measurement in a container before subtracting the difference once the powder has been transferred to a beaker. Calculating the actual concentration of salicylic acid solution Concentration of salicylic acid= no. of molesno. of moles =mass Volume Mr Original Mass of salicylic acid powder (g)| 5. 537| Mass after transfer (g)| 0. 011| Mass transferred (g)| 5. 526| =5. 526g 138. 123 =0. 400 moles = 0. 0400 1 dm3 =0. 04 moldm-3 – The powder is dissolved in aprox 100cm3 of distilled water using a magnetic stirrer and adding extra water if not all of it dissolves. – I will then rinse a 1 dm3 volumetric flask with distilled water to reduce the chance of contamination of the solution due to impurities. Then I will add the solution from the beaker, pouring it down a glass rod to avoid splashing which results in loss of some of the salicylic acid and thus a change in concentration. Similarly, I will wash the beaker with distilled water and empty it into the flask. I will then make the solution in the volumetric flask up to the 1 dm3 graduation mark with distilled water, ensure no drips will change the volume by wiping the inner-neck with absorbent paper – Having stoppered the flask securely, I will invert it several times to make sure that the molecules of salicylic acid are evenly dispersed throughout the solution to ensure a constant concentration. 2. Making up iron (III) chloride solution -This is required, along with distilled water, to dilute the salicylic acid to different concentrations to form the calibration curve.

It will allow different concentrations of the coloured complex to be formed. – To obtain accurate results, all of the salicylate ions must form complexes with the iron (III) chloride, so it must be in excess in each of my solutions. -For the colorimeter to work, the solutions must be suitably weak. I have therefore chosen 0. 01moldm-3 as the strongest concentration of salicylic acid to be used. In this I will add no distilled water and so will have to dilute the 0. 04 molar salicylic acid with iron (II) chloride in a ratio of 3:1 (iron (III) chloride: salicylic acid).

Therfore I’ll need 75 cm3 of iron (III) chloride for the 25cm3 of salicylic acid. – Given the fact that iron (III) chloride and salicylic acid form complexes in a ratio of 1:1, the number of moles of iron (III) chloride must be greater than 0. 00025 (0. 01 x 0. 025) -For the lower concentrations volumes less than 25cm3 of salicylic acid will be used, made up to 25cm3 with distilled water. Calculating the concentration of iron (III) chloride ions required so that they’re always in excess Number of moles = concentration X volume (dm3) The maximum concentration of salicylate ions is 0. 1 moldm-3. This is when 25cm3 of 0. 04 moldm-3salicylic acid solution is used. = 0. 04 X 0. 025 = 0. 001 moles So the number of moles of iron (III) chloride ions must always be in excess of 0. 001. 75 cm3 of iron (III) chloride solution is used in every concentration. So the minimum concentration of iron (III) chloride = Concentration = number of moles Volume = 0. 001 0. 075 = 0. 013 moldm-3 I have therefore chosen to make up a concentration of 0. 02 moldm-3, which will be sufficient for the ion (III) chloride ions to be in excess – The supplied iron (III) chloride solution is 1 molar.

I am going to make up 1 dm3 so as to have plenty of excess. – I will therefore pipette 20cm3 of iron (III) chloride solution into a 1 dm3 volumetric flask which has been rinsed with distilled water, and make it up to the graduation mark with distilled water. The flask will then be stoppered and inverted several times to ensure constant throughout. Planning Ethanol| Solvent for the aspirin tablets. Used instead of water due to aspirins’ low solubility in water, due to the extensive hydrogen bonding between water molecules| Hydrochloric acid| Used to make up a solution with approximately 0. moldm-3concentration, which can be used for the back titration. This is an appropriate acid as it react in a 1:1 ratio with sodium hydroxide, so similar amounts of chemicals are required| Sodium hydroxide| Used to make up a 0. 1 moldm-3 concentration solution which is used for the hydrolysis of the ester groups of the aspirin before the back titration. A low concentration is used to minimise potential error as larger volumes are therefore used. | Sodium Carbonate| Used to standardise hydrochloric acid so that its exact concentration is known.

It is appropriate as it does not degenerate. | Phenolphthalein indicator| Used in the back titration as a strong acid is reacting with a strong alkali| Methyl orange indicator| Used in the standardisation of HCL as a strong acid and a weak alkali are being used| Hydrochloric acid (II)| Used to hydrolyse the ester group of the aspirin for the colorimetry| Salicylic acid| Of approximately 0. 03M used to make up the solutions for the calibration curve for colorimetry| Iron (III) Chloride| Of about 2. moldm-3 concentration so that it can be diluted suitably, staying in excess of the salicylic acid in solution. Chosen for its property of forming a violet complex with the salicylate dianion| Distilled water| Used for all dilutions and rinsing volumetric flasks when preparing solutions, to ensure no impurity contaminates the solution affecting the results. | Branded () and generic () tablets| Required to make up the aspirin solutions, and to compare the purity of the more expensive branded tablet to a generic one. | Reasons for chemicals chosen

Equipment requirements, reasons and errors ITEM| SIZE(S)| UNCERTAINTY| REASON| 3dp electronic balance| | +/-0. 0005g| Ensures accurate mass readings of the solids required for solutions, enabling final calculations to be completed. | Pestle and Mortar| | | Allows aspirin tablets to be ground into a fine powder, aiding the dissolving process. | Colorimeter with yellow-green filter| | | Measures the absorbance of light compared to that of a sample of distilled water| Burette (x5)| 50cm3 | +/- 0. 05cm3 | For accurate measurements of volumes of solutions used in titration. Pipette (x5)| 20cm3 25cm3 50cm3 100cm3 | +/- 0. 24cm3 +/-0. 03cm3 +/-0. 08cm3 +/-0. 04cm3 | For accurate measurement of volumes of solutions required in the experiment or for making another solution. | Volumetric flask(x3)| 250cm3 1dm3 2dm3 1000 cm3 | +/-0. 30cm3 +/-0. 8cm3 +/- 0. 6 cm3 +/- 0. 08 cm3 | To achieve concentrations of exact concentrations and volumes| White tile| | | So that the end point of any titration is very clear| Conical flask (x2)| 250cm3 | +/- 0. 1 cm3| Appropriate for titrations as they reduce splashing and this loss of solution.

Their large surface area also enables colour to be easily monitored | Burette clamp and retort(2 of each)| | | Ensures stability of apparatus| Beakers| 15 small6 large| +/- 5+/- 5| Small for containing small volumes including colorimetry solutions. Large for larger volumes such as when dissolving solids in solution or combining solutions. | Cuvettes (x12)| | | Small to contain colorimetry solutions| Glass stirring rod (x4)| | | Reduces splash hen transferring solutions and aids the dissolving of solids. Plastic container bottles (x10)| | | For containing standing solutions to prevent contamination| Round-bottomed flask(x4)| | | For refluxing solutions in| Ceramic heater(x2)| | | Easy temperature control for refluxing| Condenser| | | For refluxing solutions. | Method 1- Back titration 1. Hydrolysis of aspirin solutions * Both the generic and branded aspirin solutions must be hydrolysed to salicylic acid using standardised sodium hydroxide solution * 2 moles of NaOH are consumed in the hydrolysis of 1 mole of aspirin.

To ensure hydrolysis of every aspirin molecule however, I will add a large excess of sodium hydroxide; 300cm3of NaOH per 100 cm3 of aspirin solution, giving an 6:1 molar ratio insuring excess OH ions. * Since for each type of aspirin solution I am going to have to perform several titrations to get a consistent and accurate result, I will the transfer a large volume, 100cm3, of each using pipettes into separate 500ml beakers. I will then pipette accurately into each 300cm3 of the standardised sodium hyrdroxide solution. Each mixture is then transferred to its own 500 cm3 round bottomed flask and heated under reflux for several hours, avoiding letting the solution boil. Hydrolysis is ensured by leaving it for at least 3 hours (see reflux method and diagram, page) * When the reflux is complete, stopper and invert round-bottomed flask, ensuring even distribution of molecules and so a consistent concentration. To ensure that all of the aspirin has been hydrolysed, I will add a few drops of phenolphthalein. If the mixture turns pink, it shows that NaOH is still in excess, so the aspirin has indeed all been hydrolysed. If it doesn’t turn pink, a known amount of sodium hydroxide will be added and the reflux process repeated. This will be done however many times it takes for the solution to turn pink, each time recording the amount of NaOH added) 2. Back titration * Titrate the hydrolysed solution with standard hydrochloric acid, to work out how much sodium hydroxide has reacted with acetylsalicylic acid. * (See titration method page ) Use 10cm3 portions of hydrolysed solution. * Rinse the burette and fill it with the acid solution.

Titrate it against the 10cm3 portions of hydrolysed solution (using phenolphthalein indicator, since a strong acid and strong alkali are being used, changing from pink to colourless at the end point) * I will perform rough titrations to obtain vaguely where the endpoint lies, before doing repeats for each type of aspirin until 3 readings within 0. 1 cm3 are obtained. * Calculate the umber of HCl moles used in each sample. Since the molar ratio is 1:1 this will be the same as the number of moles of NaOH in excess in the sample.

Subtracting this from the number of moles of NaOH originally in a 10cm3 sample will give the number of OH+ ions that have reacted with the ester group of the acetylsalicylic acid, and thus the number of moles of acetylsalicylic acid in the solution. This value can be used to calculate the concentration and therefore the purity of the aspirin. Method 2- Colorimetry 1. Making up the calibration solutions required for colorimetry * 10 different solutions of salicylic acid of different concentrations are required to make an accurate colorimetry calibration graph. Firstly I will carefully pipette 75cm3 of the 0. 02 moldm-3iron (III) chloride solution into each of 10 beakers. Different quantities of salicylic acid and distilled water are then required to be added to each beaker via separate pipettes. Concentration of Salicylic acid solution(moldm-3)| Volume of 0. 02 moldm-3 iron (II) chloride solution (cm3)| Volume of 0. 04 moldm-3 Salicylic acid solution (cm3)| Volume of distilled water (cm3)| 0. 000| 75. 00| 0. 00| 25. 00| 0. 001| 75. 00| 2. 50| 22. 50| . 002| 75. 00| 5. 00| 20. 00| 0. 003| 75. 00| 7. 50| 17. 50| 0. 004| 75. 00| 10. 00| 15. 00| 0. 005| 75. 00| 12. 50| 12. 50| 0. 006| 75. 00| 15. 00| 10. 00| 0. 007| 75. 00| 17. 50| 7. 50| 0. 008| 75. 00| 20. 00| 5. 00| 0. 0090. 010| 75. 0075. 00| 22. 5025. 00| 2. 500. 00| 2. Making a colorimetry calibration graph * Before any readings are taken the colorimeter must be allowed to warm up * At the advice of my teacher the colorimeter will be set at 470nm, the optimum wavelength for this type of experiment.

The cuvette must be three-quarter filled with distilled water and put into the colorimeter * This will be the benchmark (as there’s no colour present so no light should be absorbed) so I will set the absorbance to calibrate at 0. * I will then use a teat pipette to fill the next clean cuvette with my first solution (0. 01 moldm-3 ) and place it in the colorimeter to record its absorbance. Once completed, I will replace the distilled water sample in the colorimeter to reset the absorbance to 0. * This process is then carried out for the subsequent 9 concentrations.

Once complete I will be able to plot the calibration graph to show the differing absorbance for different concentrations of salicylic acid. 3. Main method of colorimetry analysis * First solutions of the generic and branded aspirin tablets are required. As in the back titration, they must both first be hydrolysed, but this time under acid conditions (hydrochloric acid) * I must hydrolyse each of my original generic and branded aspirin solutions in order to break the ester bond in the acetylsalicylic acid. * As in the alkaline hydrolysis, it is carried out by reflux (see reflux method). 0cm3 of each type of aspirin solution must be transferred to separate round-bottomed flasks via pipettes and 20cm3 of 2M hydrochloric acid added to each (the acid must be accurately pipetted as it must be in excess and the same amount for both types of aspirin) * Reflux is then carried out, and once again be left for 3 hours to ensure total hydrolysis. * Then, the same as the calibration solutions, I will add 75cm3 of 0. 02 modm3 iron (III) chloride solution into a beaker along with 25cm3 of refluxed aspirin/HCL solution, creating a total volume of 100cm3 .

To ensure that a complex forms this should then be stirred * Once again the colorimeter must warm up and be set to 470nm wavelength. Then the absorbance must once again be set to 0 by using a distilled water sample in a cuvette. * I will then ? fill another cuvette with the hydrolysed aspirin/HCL/Iron (III) chloride solution, place it in the colorimeter and record the absorbance. This will be repeated twice more, setting the absorbance to 0 in between times. * The whole process must then be repeated (reflux, colorimetry ect) for then other type of aspirin to enable comparison. 4. Analysing the results of colorimetry The colorimetry results must then be plotted on a calibration graph. This will show the readings of the differing absorbance for the generic and branded aspirin solutions. From these reading I can find their corresponding points on the calibration graph, and thus the measurements of the concentrations of the solutions and therefore the purity of the aspirin. Risk assessment Procedures * Making up and standardising solutions * Acid hydrolysis under reflux * Alkaline hydrolysis under reflux Chemical substance being used/ made| Nature of the hazard| Concentration/ Volume being used| Precautions taken| Aspirin solutions| Minimum hazard| 0. 5M, 500cm3 | Wash hands after use| Iron (III) chloride solution| Mild irritant| 0. 02 M, 1dm3| Eye protection| Salicylic acid solution| Irritant| 0. 01 M/ 0. 05M| Eye protection| Sodium hydroxide solution| 0. 1M-irritant2M- corrosive| 0. 1 M| Eye protectionGloves for high concentration| Ethanol| Highly flammable| absolute| Eye protection| Hydrochloric acid| 1M-irritant2M- corrosive| 0. 1M/ 2M| Eye protection| Sodium carbonate solution| Minimum | 0. 1 M, 500cm3| Eye protection| Non-chemical hazards ;amp; precautions taken * Remove bags from walkways and put stools under the desks to avoid tripping * Tie long hair back Wear a lab-coat to protect clothing * Safety glasses at all times * Clear any spillages/ breakages immediately Waste disposal * Most can be disposed of in sinks, any organic waste (none here) must go into organic waste bottles. Results i) back titration 1. Making up branded aspirin solution Mass of 5 tablets: 2. 155 g Mass of 1 tablet: 0. 431 g Initial tared mass of branded aspirin tablets: 5. 974 g Tared mass of receptacle and residual tablet powder: 0. 125 g Mass of tablet used: 5. 849 g Total volume of solution (made up with ethanol): 500 cm3 2. Making up generic aspirin solution

Mass of 5 tablets: 3. 008 g Mass of 1 tablet: 0. 6016 g Initial tared mass of branded aspirin tablets: 9. 041 g Tared mass of receptacle and residual tablet powder: 0. 094 g Mass of tablet used: 8. 947 g Total volume of solution (made up with ethanol): 500 cm3 3. Making up sodium carbonate solution of known concentration Initial tared mass of sodium carbonate: 5. 372 g Tared mass of receptacle and residual sodium carbonate: 0. 028 g Mass of sodium carbonate used: 5. 344 g Total volume of solution (made up with water): 500 cm3 4. Standardising HCl for the back titration

Factors kept constant: eye level when making recordings from the burette Concentration of sodium carbonate: 0. 1008 moldm-3 Volume of sodium carbonate used: 20 cm3 Initial burette reading HCl (cm3 )| End burette reading HCl (cm3 )| Volume of HCl used (cm3 )| 0. 200| 41. 400| 41. 200| 0. 200| 41. 700| 41. 200| 0. 300| 41. 400| 41. 100| Average volume of HCl used (cm3 )| | 41. 167| 5. Standardising sodium hydroxide for the back titration Factors kept constant: eye level when making recordings from the burette Concentration of HCl: 0. 098 moldm-3 Volume of sodium NaOH used: 20 cm3

Initial burette reading HCl (cm3 )| End burette reading HCl (cm3 )| Volume of HCl used (cm3 )| 0. 200| 19. 600| 19. 400| 0. 500| 19. 800| 19. 300| 0. 500| 19. 900| 19. 400| Average volume of HCl used (cm3 )| | 19. 367| 6. Back titration of branded aspirin solution Factors kept constant: eye level when making recordings from the burette Volume of NaOH- aspirin solution used: Concentration of HCl: Initial burette reading HCl (cm3 )| End burette reading HCl (cm3 )| Volume of HCl used (cm3 )| 12. 100| 17. 200| 5. 100| 17. 300| 22. 500| 5. 200| 22. 500| 27. 700| 5. 200| Average volume of HCl used (cm3 )| | 5. 167|

Initial burette reading HCl (cm3 )| End burette reading HCl (cm3 )| Volume of HCl used (cm3 )| 4. 900| 9. 600| 4. 700| 8. 500| 13. 300| 4. 800| 12. 100| 16. 900| 4. 800| Average volume of HCl used (cm3 )| | 4. 767 (3dp)| 1. Back titration of generic aspirin solution Factors kept constant: eye level when making recordings from the burette Volume of NaOH- aspirin solution used: 10 cm3 Concentration of HCl: 0. 098 moldm-3 ii) colorimetry 1. Table showing volumes of solutions used to prepare calibration solutions Factors kept constant: total volume of each solution = 100 cm3 Volume of iron (III) chloride = 75cm3 Having made up solutions of the shown constitutions, I found the colorimeter I was using to be too sensitive and absorbance readings came out all the same, at 2. 5, which is the maximum reading. – I therefore decided to dilute each of my solutions tenfold – For each solution, I used a 10cm3 pipette to transfer 10cm3 to a 100 cm3 flask. Concentration of Salicylic acid solution (moldm-3)| Volume of 0. 02 moldm-3 iron (II) chloride solution (cm3 )| Volume of 0. 04 moldm-3 Salicylic acid solution (cm3 )| Volume of distilled water (cm3 )| 0. 000| 75. 00| 0. 00| 25. 00| 0. 001| 75. 00| 2. 50| 22. 50| . 002| 75. 00| 5. 00| 20. 00| 0. 003| 75. 00| 7. 50| 17. 50| 0. 004| 75. 00| 10. 00| 15. 00| 0. 005| 75. 00| 12. 50| 12. 50| 0. 006| 75. 00| 15. 00| 10. 00| 0. 007| 75. 00| 17. 50| 7. 50| 0. 008| 75. 00| 20. 00| 5. 00| 0. 0090. 010| 75. 0075. 00| 22. 5025. 00| 2. 500. 00| – I then added distilled water, using a teat pipette when it neared the graduation mark, to bring the solution up to where the base of the meniscus touched the graduation mark. – I then wiped the inside of the neck of the volumetric flask with absorbent paper to prevent excess distilled water dripping into the solution, altering its concentration. I then securely stoppered the flask and inverted it several times to ensure a constant concentration throughout the solution. -This same method of dilution was performed in the actual aspirin solutions, in order that they would read on the scale of the colorimeter, and to be consistent with the salicylic acid solutions. 2. Table showing absorbance by calibration solutions Factors kept constant: Wavelength of light = 470 nm Concentration of Salicylic acid solution (moldm-3)| Absorbance | 0. 000| 0. 000| 0. 0001| 0. 201| 0. 0002| 0. 291| 0. 0003| 0. 375| 0. 0004| 0. 522| 0. 0005| 0. 631| 0. 0006| 0. 772| 0. 0007| 0. 55| 0. 0008| 1. 020| 0. 0009| 1. 144| 0. 0010| 1. 282| 3. Table showing absorbance by aspirin solutions | Repeat 1| Repeat 2| Repeat 3| Average absorbance| Branded aspirin| 0. 291| 0. 314| 0. 285| 0. 297| Generic aspirin| 0. 354| 0. 371| 0. 412| 0. 379| 4. Calibration curve Concentration of salicylic acid solution – I plotted this concentration-absorbance graph using Microsoft excel as graph drawing software. I also plotted a least squares regression line as a line of best fit for the data. R2 (how well the data fits, out of 1) = 0. 9957 Equation of the line of best fit: y = 1252. 6x +0. 0276 Analysis of results Calculating the exact concentrations of: -sodium carbonate solution -hydrochloric acid solution – sodium hydroxide solution -aspirin solution (based on packet-labelled strength) * Calculating the purity of aspirin from the back titration results * Calculating the purity of aspirin from the results from the colorimetry Summary of the preparation and standardisation of solutions Solution| Calculation page reference| Concentration (M)| Standardised hydrochloric acid| Page 13| 0. 098| Standardised sodium hydroxide| Page 11| 0. 095| Standardised sodium carbonate| Page 14| 0. 101| Branded aspirin*| Page 9| 0. 490| Generic aspirin*| Page 9| 0. 0495| *concentrations based on the labelled strengths of the tablets on their packaging Back titration results Branded aspirin tablet: Finding the number of moles of aspirin present in each sample of 10cm3 hydrolysed solution Step 1: volume of HCL required in back titration | Volume of HCL required to reach endpoint (cm3 )| Titre 1| 5. 1| Titre 2| 5. 2| Titre 3Average| 5. 25. 167| Step 2: moles of HCL used No. of moles= concentration X volume =0. 098moldm-3 X 0. 005167 dm3 =5. 06366 X 10-4 Step 3: moles of NaOH in excess in 10cm3 sample of aspirin solution

Since HCl and NaOH react in a 1:1 ratio, the moles of sodium hydroxide available to react with the HCl is indicative of the number of moles in excess and is therefore the same as the no. of moles of HCl used in the titration. Moles of NaOH= 5. 06366 X 10-4 Step 4: total number of moles of NaOH in 10cm3 sample – First I need to know the number of moles of NaOH in my entire aspirin solution Number of moles= volume of NaOH used X concentration of standardised NaOH = 0. 3dm3 X 0. 095 =0. 0285 moles – Now I need to calculate the concentration of NaOH in the aspirin solution, as it has now been diluted due to the aspirin solution

I used 300cm3 of NaOH and 100 cm3 of aspirin solution making a total volume of 400cm3 Concentration of NaOH = number of moles Volume = 0. 0285 0. 4dm3 =0. 071moldm-3 – Now I can work out how many moles of NaOH were in each of my 10cm3 samples used in the titration Number of moles = concentration X volume = 0. 071 moldm-3 X 0. 01dm3 =7. 125 X 10-4 Step 5: moles of NaOH used in hydrolysis = total moles of NaOH in 10cm3 sample – moles of NaOH in excess = 7. 125 X 10-4 – 5. 064 X 10-4 = 2. 061 X 10-4 Step 6: moles of aspirin in 10 cm3 sample

Knowing that aspirin and NaOH react in a ratio of 1:2 we can calculate the moles of aspirin from the moles of NaOH used in hydrolysis = 1. 031 X 10-4 However, to obtain the % purity of the branded aspirin, this value for the measured amount of aspirin in a 10cm3 sample must be divided by the theoretical quantity of aspirin in a 10cm3 sample. Branded solution: calculating the theoretical amount of aspirin in a 10cm3 sample based on stated tablet strength (according to the packet) Step 7: theoretical concentration of aspirin and NaOH in solution prior to hydrolysis (based on labelled strength of tablets) I found the theoretical concentration of branded aspirin solution to be 0. 0490moldm-3, however it is then diluted by combining it with NaOH (100cm3 of aspirin solution was combined with 300 cm3 of NaOH making a total solution of 400 cm3 ) Moles of aspirin in 100cm3 = concentration X volume =0. 0490 moldm-3 X 0. 1dm3 = 4. 9 X 10-3 moles new concentration of aspirin= number of moles total volume =4. 9 X 10-3 0. 4 dm3 = 0. 01225 moldm-3 =0. 0123moldm-3 Step 8: From the theoretical no. of moles of aspirin in my solution prior to hydrolysis, I cant calculate the number of moles in each 10cm3 sample No. f moles in 10cm3 = concentration X volume =0. 0123 moldm-3 X 0. 01 dm3 =1. 23 X 10-4 Step 9: Finally, to find the % purity of the branded aspirin, the results for steps 6 and 8 must be compared: This compares the values for the experimentally calculated quantity of aspirin in a 10cm3 sample (step 6) with the theoretical value (step 8). A ratio of theses two will give the % purity. Result of step 6 X 100 = measured % purity Result of step 8 = 1. 031 X 10-4 X 100 1. 23 X 10-4 Branded aspirin tablets =83. 821 % purity of acetylsalicylic acid Generic aspirin tablet Step 1: volume of HCL required in back titration Volume of HCL required to reach endpoint (cm3 )| Titre 1| 4. 700| Titre 2| 4. 800| Titre 3Average| 4. 8004. 767| Step 2: moles of HCL used No. of moles= concentration X volume =0. 098moldm-3 X 0. 004767 dm3 =4. 672 X 10-4 Step 3: moles of NaOH in excess in 10cm3 sample of aspirin solution Since HCl and NaOH react in a 1:1 ratio, the moles of sodium hydroxide available to react with the HCl is indicative of the number of moles in excess and is therefore the same as the no. of moles of HCl used in the titration. Moles of NaOH= 4. 672 X 10-4 Step 4: total number of moles of NaOH in 10cm3 sample First I need to know the number of moles of NaOH in my entire aspirin solution Number of moles= volume of NaOH used X concentration of standardised NaOH = 0. 3dm3 X 0. 095 =0. 0285 moles – Now I need to calculate the concentration of NaOH in the aspirin solution, as it has now been diluted due to the aspirin solution I used 300cm3 of NaOH and 100 cm3 of aspirin solution making a total volume of 400cm3 Concentration of NaOH = number of moles Volume = 0. 0285 0. 4dm3 =0. 071moldm-3 – Now I can work out how many moles of NaOH were in each of my 10cm3 samples used in the titration Number of moles = concentration X volume 0. 071 moldm-3 X 0. 01dm3 =7. 125 X 10-4 Step 5: moles of NaOH used in hydrolysis = Total moles of NaOH in 10cm3 sample – moles of NaOH in excess = 7. 125 X 10-4 – 4. 672 X 10-4 = 2. 453 X 10-4 Step 6: moles of aspirin in 10 cm3 sample Knowing that aspirin and NaOH react in a ratio of 1:2 we can calculate the moles of aspirin from the moles of NaOH used in hydrolysis = 1. 2265 X 10-4 =1. 227 X 10-4 However, to obtain the % purity of the generic aspirin, this value for the measured amount of aspirin in a 10cm3 sample must be divided by the theoretical quantity of aspirin in a 10cm3 sample.

Generic solution: calculating the theoretical amount of aspirin in a 10cm3 sample based on stated tablet strength (according to the packet) Step 7: theoretical concentration of aspirin and NaOH in solution prior to hydrolysis (based on labelled strength of tablets) – I found the theoretical concentration of generic aspirin solution to be 0. 0495moldm-3, however it is then diluted by combining it with NaOH (100cm3 of aspirin solution was combined with 300 cm3 of NaOH making a total solution of 400 cm3 ) Moles of aspirin in 100cm3 = concentration X volume 0. 0495 moldm-3 X 0. 1dm3 = 4. 95 X 10-3 moles New concentration of aspirin= number of moles total volume =4. 95 X 10-3 0. 4 dm3 = 0. 012375 moldm-3 =0. 0124moldm-3 Step 8: From the theoretical no. of moles of aspirin in my solution prior to hydrolysis, I cant calculate the number of moles in each 10cm3 sample No. of moles in 10cm3 = concentration X volume =0. 0124 moldm-3 X 0. 01 dm3 =1. 24 X 10-4 Step 9: Finally, to find the % purity of the generic aspirin, the results for steps 6 and 8 must be compared:

This compares the values for the experimentally calculated quantity of aspirin in a 10cm3 sample (step 6) with the theoretical value (step 8). A ratio of theses two will give the % purity. Result of step 6 X 100 = measured % purity Result of step 8 = 1. 227 X 10-4 X 100 1. 24 X 10-4 generic aspirin tablet = 98. 952 % pure acetylsalicylic acid Colorimetry results Step 1: Plotting the calibration curve The mean absorbance of light is plotted against the concentration of salicylic acid, to show the relationship between the two. (see page 28 for graph) Step 2 : Calculate the mean absorbance for each aspirin- iron (III) chloride solution calculations see p 29) To gain the most reliable and precise representation of the whole experiment I am using just the mean value for absorbance for each aspirin solution. One value is easier to compare against other results. Branded aspirin absorbance = 0. 297| Generic aspirin absorbance = 0. 379| Step 3: Use the calibration graph to obtain the initial concentration of aspirin. The ratio of salicylic acid to aspirin in hydrolysis is 1:1. By using the equation for the regression line, I can calculate the concentration of aspirin initially in the each of the two colorimetry samples. Y = 1252. 6x +0. 0276 -Absorbance for branded aspirin: 0. 97 0. 297 = 1252. 6x + 0. 0276 where x = concentration of salicylic acid Concentration of salicylic acid and thus aspirin in sample: 2. 15 x 10-4 mol dm-3 -Absorbance for generic aspirin: 0. 379 0. 379 =1252. 6x +0. 0276 Concentration of salicylic acid and thus aspirin in sample: 2. 81 x 10-4 mol dm-3 Step 4: Calculate the theoretical concentration of aspirin in each hydrolysed aspirin- iron (III) chloride solution before hydrolysis (theoretical: based on box-labelled strength) Branded: theoretical concentration = 0. 0490 M However, this solution was then diluted by the following factors in the acid hydrolysis and colorimetry: 50 cm3 of aspirin solution was combined with 20 cm3 of HCl in acid hydrolysis dilution ratio : 50/70 -25 cm3 of the hydrolysed solution was then combined with 75 cm3 of iron (III) chloride solution for the colorimetry dilution ratio : 25/100 – This was then diluted tenfold in the same way as the salicylic acid solutions dilution ratio: 10/100 So the theoretical original concentration of branded aspirin in the colorimetry sample was: 0. 0490 x (50/70) x (25/100) x (10/100) = 0. 000875 M Generic: theoretical concentration = 0. 0495 M This aspirin solution was diluted in the same ratios as the branded aspirin solution

So. The the theoretical original concentration of generic aspirin in the colorimetry sample was: 0. 0495 x (50/70) x (25/100) x (10/100) = 0. 000884 Step 5: Calculate the measured purity of the aspirin by the ratio step 3:step 4 Branded: result of step 3 x 100 Result of step 4 = 2. 15 x 10-4 x 100 0. 000875 = 24. 571 % purity (3dp) Generic: result of step 3 x 100 Result of step 4 = 2. 81 x 10-4 x 100 0. 000884 = 31. 787 % purity (3dp) | % purity in Back titration| % purity in Colorimety| Generic aspirin solution| 98. 952| 31. 787| Branded aspirin solution| 83. 821| 24. 571|

This table enables my intended comparison of not only the stated purity of aspirin with the measured purity, but also the two methods of composition analysis. Since I measured the purity of aspirin relative to the stated amount of aspirin per tablet (according to packet labelled strength), the measured purity should be 100 %. 100 % is therefore my expected value for generic and branded aspirin, in both procedures. Since, as shown above, my measured purities were all below this value, uncertainties and shortcomings in my procedures must be considered (see evaluation) to explain this difference.

The UK’s very strict laws on the quality of pharmaceuticals means that I can conclude that one of these two factors must be responsible for the difference, the only exception being If the tablets were very stale, in which case you might expect a purity marginally lower than the expected value of 100 %. I can conclude now that the back titration method of analysis was more accurate than colorimetry, as it produced % purities far closer to the expected value. The % purity for generic aspirin of 98. 952 % is very close to the expected value. However, once again uncertainties must be used to xplain the lower % purity for the branded aspirin. In my evaluation section I will discuss potential sources of uncertainty for both methods of analysis, however, any difference in uncertainty alone is not enough to explain such a large difference in accuracy for the two methods. I must therefore find the flaws in experimental procedure in the colorimetry. Evaluation In order to enable valid comparison of the two methods of composition analysis, I must include and consider any uncertainties which could affect the results. % uncertainty = measured uncertainty x 100

Amount of substance I must therefore add up the % uncertainties of each part of the two methods to include them in the overall calculated purity result. Back titration uncertainties Branded solution Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| Electronic balance| +/- 0. 0005g x 2| 5. 849g| 0. 0855| Volumetric flask (500cm3 )| +/- 0. 2cm3 (class B)| 500cm3 | 0. 0400| Total % uncertainty| | | 0. 1255| Generic solution Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| Electronic balance| +/- 0. 0005g x 2| 8. 47g| 0. 00559| Volumetric flask (500cm3 )| +/- 0. 2cm3 (class B)| 500cm3 | 0. 0400| Total % uncertainty| | | 0. 04559| The uncertainties for the electronic balance and multiplied by two to account for the mass of aspirin being more accurately measured by difference; this therefore take into account the fact that two weightings have taken place. Uncertainties making up 0. 1 molar sodium carbonate solution Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| Electronic balance| +/- 0. 0005g (x 2)*| 5. 344g| 0. 00936| Volumetric (500cm3 )| +/- 0. cm3 (class B)| 500cm3 | 0. 0400| Total % uncertainty| | | 0. 04936| So my standardised solution of sodium hydroxide, of concentration 0. 1008moldm-3, has the uncertainty: +/- 0. 04936 Uncertainties when standardising HCl solution Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 20 cm3 pipette| +/- 0. 024 cm3 | 20 cm3 | 0. 1200| 50 cm3 burette| +/- 0. 05 cm3 (x2)**| 41. 167| 0. 1210| Sodium carbonate| | | 0. 04936| Total % uncertainty0. 29036 **burette measurements are done by difference, hence uncertainty times by 2. So the standardised 0. 98 molar HCl has uncertainty: +/-0. 29036 Uncertainties when standardising sodium hydroxide Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 20 cm3 pipette| +/- 0. 024 cm3 | 20cm3 | 0. 1200| 50 cm3 burette| +/- 0. 05 cm3 (x2)**| 19. 367cm3 | 0. 2580| Standardised HCl| | | 0. 29036| Total % uncertainty:0. 66836 For the standardised 0. 095 molar NaOH has the % uncertainty : +/-0. 66836 Uncertainties during the back titration procedure Branded aspirin Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 10 cm3 pipette| +/-0. cm3 | 10 cm3 | 1. 0000| 50 cm3 burette| +/- 0. 05 cm3 (x2)**| 5. 167 cm3 | 0. 9680| Standardised HCl| | | 0. 29036| Total % uncertainty in measured concentration| | | 2. 25836| OVERALL UNCERTAINTY IN PURITY (%)| | | +/- 2. 25836 %| Total measured % purity of branded aspirin tablets in back titration = 83. 821 +/- 2. 258 % This gives a range of % purity: 86. 079 – 81. 563 % Generic aspirin Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 10 cm3 pipette| +/-0. 1 cm3 | 10 cm3 | 1. 0000| 50 cm3 burette| +/- 0. 05 cm3 (x2)**| 4. 767 cm3 | 1. 4888| Standardised HCl| | | 0. 29036| Total % uncertainty in measured concentration| | | 2. 33924| OVERALL UNCERTAINTY IN PURITY (%)| | | +/- 2. 33924 %| Total measured % purity of generic aspirin tablets in back titration = 98. 952 +/- 2. 339 % This gives a range of % purity: 101. 291 – 96. 613 % Comments on uncertainties During the back titration, the use of burettes with only small volumes in them accounted for the greatest uncertainties. There was uncertainty due to having to read from the scale. The burette uncertainties are also doubled as the titres are measured by difference.

This affects the overall measurement for aspirin purity as it affects the volume of HCl actually used, and thus the number of moles of reacted NaOH in hydrolysis. To reduce this uncertainty in the future I would dilute the HCl concentration by a known amount, to increase the titration volume and thus reduce measurement errors. Also, I could make a greater excess of sodium hydroxide, so that a greater volume of hydrochloric acid is required to neutralise it. Alternatively, I could use an even smaller volume of aspirin solution to analyse, however this volume may be insufficient to produce reliable results.

It is also possible that human error could have increased the uncertainty in this step (the actual back titration) more than in others. This is because I performed the back titration at the end of the day, having done a full day of standardising solutions and refluxing, so tiredness could have had an impact on accuracy. In making up the aspirin solutions uncertainty was much smaller. For this reason I believe using class A rather than the class B volumetric flasks used would have produced a negligible difference.

A similarly minimal difference in uncertainty would be achieved by using an electronic balance with a higher degree of accuracy, again considering the overall uncertainty was so small anyway for this step, and that the presence of many people in the lab meant that the accuracy of the balance was dependant on no one else tampering with it anyway. Human error should not have contributed to uncertainty in this step as I made up the aspirin solutions at the very beginning of the day. I therefore conclude that, for this step, there were no changes in method which could have significantly reduced uncertainty in the given lab conditions.

In the making up of all solutions, including aspirin solutions, hydrochloric acid and sodium carbonate, large volumes were made up (500 cm3 ) which I know reduces the uncertainty in the volumetric, thus reducing percentage error. 10 In the standardisation of HCl and NaOH, the biggest uncertainties once again came from the use of small volumes in 50 cm3 burettes. The use of one standardised solution to standardise another (eg sodium carbonate for HCl and then standardised HCl for sodium hydroxide) had a big impact, as uncertainties previously calculated must be added on to the extra ncertainty of each new concentration calculation. The uncertainty could have been reduced by the use of a primary standard (eg oxalic acid) for the standardisation of sodium hydroxide, rather than HCl10. For all of theses titrations, I could once again have used a reduced concentration of each chemical in the burette, making larger volumes and reducing uncertainty. In all I believe that human error (inaccurate use of apparatus) made little contribution to the measurement uncertainties compared to some of the other sources.

The first things I would change would be to use more dilute solutions and larger volumes of them to reduce the effect of percentage uncertainty, and would in some cases choose class A burettes rather than class B. General comments on my back titration results ;amp; other uncertainties In my back titration method, I obtained purity values of 101. 291 – 96. 613 % and 86. 079 – 81. 563 % for generic and branded aspirin tablets respectively. Having measured purity in respect to the stated amount of aspirin per tablet, my expected value was 100 %.

My measured purity for generic aspirin is therefore as expected, as the equipment uncertainty in the procedure accounts for the small margin that it is off 100 %, as shown in the range of purities for generic aspirin above. From these results, I could conclude that generic aspirin is more pure than branded aspirin, as the purity ranges do not overlap. However, the value of 83. 821 % purity for the branded aspirin, even taking into account equipment uncertainty and considering inevitable human error in the procedure, this value is still significantly less than its expected value of 100 %.

I must therefore consider possible explanations for this difference. One area in which error could have occurred is the alkaline hydrolysis prior to the back titration. It is possible that the acetylsalicylic acid of the branded aspirin solution was not fully hydrolysed, potentially due to not being at optimum conditions for the hydrolysis reaction. In future, I could reflux the mixture for further time and with a higher proportion of alkali in comparison to the aspirin. However I think that this explanation is unlikely because I refluxed for several hours as it was and calculated the amount of alkali added to be in plenty of excess.

Also, the reflux of the generic and branded aspirin solutions were performed identically, and the generic aspirin solution showed no signs of incomplete hydrolysis as it gave a measured purity of very close to 100 %. Another area that requires more research is the possibility of impurities or other ingredients in the aspirin tablets interfering with the hydrolysis. The generic aspirin variety, according to the leaflet within the packet, contains potato starch, talc and lactose. Similarly, the branded variety contains a form if starch; maize starch.

Other additives such as microcrystalline cellulose, caffeine or quinine sulphate in the branded tablets could have lead rise to the incomplete hydrolysis of the branded aspirin solution, and thus the difference in purity between the generic and branded solutions. Further research into the potential effects of these impurities on the hydrolysis reaction and other stages of my procedure is required to conclusively identify the cause of the difference between expected and actual measure purity for the branded aspirin.

The human source of uncertainty in this procedure is the judgement of the endpoint of the neutralisation in the back titration. This is a likely source of error, as the colour change is immediate and it is hard to shut of the tap of the burette at exactly the right time. Also, determining the moment of colour change was challenging for the branded solution in particular, as an impurity in the branded tablets caused the hydrolysed solution to go a pale brown colour.

Further research into the impurity which could have caused this premature colour change is required to be able to conclude that this is the uncertainty causing the lower than expected purity for the branded tablets. As discussed above, due to the measured purity for branded aspirin being considerably less than the expected value, and thus the presence of an error in procedure greater than simple equipment uncertainties, I cannot conclude which type of aspirin is purer. The reliability of the experiment must be mproved before conclusions can be drawn, by investigation of non-measurement errors and further repeats of the experiment. Uncertainties in colorimetry method of analysis Making up Salicylic acid Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 1 dm3 volumetric flask| +/- 0. 8 cm3 | 1 dm3 | 0. 0800| Electronic balance| +/- 0. 0005 (x2)*| 5. 526 g| 0. 01810| Total % uncertainty| | | 0. 261%| Once again, the uncertainty of the electronic balance is multiplied by 2 due to the salicylic acid powder being measured by difference. Making up iron (III) chloride solution

Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 1 dm3 volumetric flask| +/- 0. 8 cm3 | 1 dm3 | 0. 0800| 20 cm3 pipette| +/- 0. 024 cm3 | 20cm3 | 0. 1200| Total % uncertainty| | | 0. 20%| This total uncertainty of 0. 2 % has the same impact on all of the calibration solutions Making up calibration curve solutions Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 100 cm3 pipette| +/- 0. 04 cm3 | 75 cm3 | 0. 0533| 25 cm3 pipette (salicylic acid)25cm3 pipette (water)| +/- 0. 03 cm3 +/- 0. 03 cm3 | 2. 5- 25 cm3 2. – 25 cm3 | Total uncertainty between: 1. 2 and 0. 12 %| Iron (III) chloride| | | 0. 2000| Total % uncertainty| | | +/- 1. 4533- 1. 5733 %| The uncertainties for the 25cm3 pipettes used are most significant when making up the strongest and weakest concentrations of salicylic acid, as +/-0. 03 cm3 uncertainty is significant relative to a very small volume like 2. 5 cm3 . Since these absolute uncertainties are small, I chose not to plot error bars on my calibration curve given the scale of the graph. On my calibration curve I chose to plot a least squares regression line.

Since I could have plotted a different line, this introduces more uncertainty in the actual reading of concentrations from the graph (by use of the formula created). This uncertainty in the use of the graph to obtain concentrations of the aspirin solutions could in fact be more significant than the measured uncertainty. Uncertainty for aspirin solutions Branded Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 50 cm3 pipette (x2)| +/- 0. 08 cm3 | 50 cm3 | 0. 3200| 25 cm3 pipette (x2)| +/- 0. 03 cm3 | 25 cm3 | 0. 2400| 0 cm3 pipette| +/- 0. 24 cm3 | 20 cm3 | 1. 2000| 10 cm3 pipette (used in dilution)| +/- 0. 02 cm3 | 10 cm3 | 0. 2000| 100cm3 volumetric flask (used in dilution)| +/- 0. 03 cm3 | 100 cm3 | 0. 0300| Total % uncertainty| | | +/- 1. 99%| Using these uncertainties, the measured purity of branded aspirin from the colorimetry method is 24. 571 +/- 1. 99 % This gives a range of purities of the branded aspirin from 26. 561 – 22. 581 % Generic Equipment| Uncertainty in measurement| Amount of material| % Uncertainty (+/-)| 50 cm3 pipette (x2)| +/- 0. 8 cm3 | 50 cm3 | 0. 3200| 25 cm3 pipette (x2)| +/- 0. 03 cm3 | 25 cm3 | 0. 2400| 20 cm3 pipette| +/- 0. 24 cm3 | 20 cm3 | 1. 2000| 10 cm3 pipette (used in dilution)| +/- 0. 02 cm3 | 10 cm3 | 0. 2000| 100cm3 volumetric flask (used in dilution)| +/- 0. 08 cm3 | 100 cm3 | 0. 0300| Total % uncertainty| | | +/- 1. 99%| Using these uncertainties, the measured purity of generic aspirin using the colorimetry method is 31. 787 +/- 1. 99 % This gives a range of purities from 33. 777 – 29. 797 % Comments on results of colorimetry and uncertainties

Even taking into account these uncertainties, the measured purities obtained from the colorimetry are much lower than those from the back titration, and are far from the expected value of 100%, suggesting a systematic error in the method. Considering potential sources for this error, my first thought was the sensitivity of the colorimeter. I already had to dilute all of my solutions of salicylic acid tenfold to account for such sensitivity, and so have now investigated whether 470 nm was in fact the appropriate wavelength setting. For other similar experiments I researched and found the optimum wavelength to be 530nm 11.

From this source I also read about the use of buffers to improve the reliability if such colorimetric analysis. This is because, if the solution if too acidic, the purple colour of the complex (salicylic acid with iron (III) chloride) does not form, whereas if the pH is too high, the solution becomes cloudy due to the hydrolysis of the hydrated iron (III) ion to iron (III) hydroxide, affecting the absorbance reading. Therefore to obtain more accurate and reliable results, I would suggest a repeat of the colorimetry using a buffer and the more appropriate, higher wavelength of 530 nm on the colorimeter.

These amendments should solve the problem of obtaining the desired purple colour, and would in doing so change the curvature of the calibration curve. A change in this line would lead to different readings for the purity of the aspirin solutions. Another potential error in my procedure, leading rise to such low purities, is the incomplete hydrolysis of the aspirin, which would in turn have caused a lower concentration of salicylic acid and thus the low absorbance readings for the experimental solutions. It is possible that the conditions (eg pH/time) for the acid hydrolysis were not optimum, or that it is simply less