Analyse the purity of my aspirin samples (aspirin and recrystallized aspirin) compared to commercial aspirin by use of analysing melting points, TLC plates and colorimetry tests Investigate how pH effects the rate of hydrolysis of aspirin Background theory Making my own sample of Salicylic acid (to then be converted into Aspirin) Salicylic acid (2-Hydroxybenzoic acid) can be made by hydrolysing methyl 2-hydroxybenzoate. Oil of wintergreen is 98% methyl 2-hydroxybenzoate and is what I used as my starting chemical. The structure of methyl-2-hydroxybenzoate is shown on the right, and is an ester as it has the functional group –COO-.
Hydrolysing esters involves splitting them into carboxylic acids or their salts and alcohols by adding water, dilute acid or dilute alkali. To hydrolysed my oil of wintergreen I first reacted it with a weak alkali, aqueous sodium hydroxide, and heated it under reflux to produce sodium 2-hydroxybenzoate. I then converted this into 2-hydroxybenzoic acid by adding hydrochloric acid. Technically, hydrolysis is a reaction with water, however, the exact thing happens when acids/alkalis hydrolyse esters, the reaction is just quicker as the acid/alkali acts as a catalyst.
To hydrolyse the methyl-2-hydroxybenzoate it is heated under reflux with the dilute alkali sodium hydroxide. First we hydrolyse the methyl-2-hydroxybenzoate using aqueous sodium hydroxide: The sodium salt, sodium 2-hydroxybenzoate, is formed rather than the carboxylic acid itself (which is what we want). To get the acid (Salicylic acid) we now react the sodium 2-hydroxybenzoate with an excess of conc. Hydrochloric acid. Doing this floods the mixture with hydrogen ions which are picked up by the hydroxybenzoate ions present in the sodium salt to make 2-Hydroxybenzoic acid (salicylic acid) as shown below.
The 2-hydroxybenzoic acid can now be filtered off and we have our sample of salicylic acid. Converting Salicylic acid into aspirin Aspirin is an acetate ester of the aromatic alcohol, salicylic acid. Salicylic acid is converted into aspirin by esterification. Esterification is the reaction of a carboxyl group and an -OH group of an alcohol or phenol in the presence of an acid catalyst to form a carboxylate ester and is reversible. The reaction reaches equilibrium during refluxing. In the synthesis of aspirin the -OH group is the phenolic -OH group attached to ring of the salicylic acid as shown in the equation below.
The acetyl group ( -COCH3) comes from ethanoic anhydride (acetic anhydride). The reaction in my first method was catalyzed by phosphoric acid, H3PO4, in my second method I used sulphuric (VI) acid, H2SO4. Phosphoric acid is a weaker acid than sulphuric acids. The reaction takes place easily in the acidic solution but the product is formed as part of a mixture containing several other compounds. Therefore the product is formed and then separated from impurities. The equation below shows the reaction as I have described.
It is important to note that this is a reversible reaction, so the salicylic acid and ethanoic anhydride reactants and aspirin product are in equilibrium. This means that once the reaction reaches equilibrium, there could be a large amount of starting material remaining, resulting in a poor yield of aspirin. To resolve this problem we use Le Chatelier’s principle, which predicts that we can drive the equilibrium to the right (to the products) by having one of the reactants in excess. This basically puts pressure on the left side, which allows the equilibrium to move to the right. The acetic anhydride or salicylic acid can be used in excess.
The choice is based on cost, availability and/or ease of purification at the end of the reaction. In my reaction I added an excess acetic anhydride. A 1:1 mixture of acetic anhydride and Salicylic acid will yield an equilibrium mixture that is about 70% aspirin. This means that if the aspirin were isolated from this mixture, in perfect conditions a 70% yield would be obtained. Using Le Chatelier’s principle again, if a ratio of 3:1 or 1:3 of acetic anhydride: salicylic acid were used, the equilibrium would be driven towards the aspirin product and in a perfect system would result in a 90% yield.
In practice often a 10:1 or 1:10 ratio is used, which results in even higher yields of product. The disadvantage of this method is that most of the excess reactant is left unreacted, requiring extra steps to remove it that I will address below. Sulfuric acid (H2SO4) is used as a catalyst for the reaction and accelerates the rate at which the aspirin is formed. Since a catalyst is not consumed during the course of a reaction, I needed to use only a small amount of sulfuric acid in order for it to be effective. Heating is another way to increase the reaction rate.
Heating the mixture increases the rate of reaction as the heat energy is converted into kinetic energy of the molecules. The more kinetic energy the molecules have (the more heat) the larger the number of collisions with sufficient energy to overcome to activation energy barrier and result in successful collisions. This increases the rate of reaction as more successful collisions are able to take place over a certain period of time meaning more product is formed over a certain period of time. I heated the reaction using reflux apparatus (see method).
Refluxing is a process where the reaction mixture is heated to its boiling point, with a condenser attached, so that when the liquid vaporizes it rises up the condenser but condenses and falls back down into the reaction flask. This process allows the reaction to be heated over a period of time, without evaporating away the solvent or reactants. After esterification has taken place cold distilled water is added to the reaction mixture which causes the precipitation of the aspirin and removes the impurities in the solution as it still contains unreacted sulphuric acid, excess acetic anhydride and acetic acid.
The water reacts with acetic anhydride to form 2 molecules of acetic acid, which is water soluble. The acetic acid dissolves in water because both the acetic and water molecules are polar. Acetic acid has a partial negative charge on the oxygen atom and a partial positive charge on the hydrogen. Ice cold distilled water is used when collecting the aspirin instead of room temperature distilled water because aspirin is insoluble in cold water and therefore the aspirin is not dissolved. Whereas acetic acid and sulphuric acid are water soluble in the ice cold water. It is important that the reaction takes place in anhydrous conditions.
This is because acetic acid is a weak acid and therefore when H2O is added it does not fully dissociate in solution. However, some of the acetic acid and the water react and form hydronium ions, H3O+. This causes the reverse reaction to be favored as the equilibrium of the reaction is more favored towards the reactants than the products. We do not want this because we are trying synthesize the product aspirin. I tested my collected aspirin for it’s purity using FeCl3 (aq). The Fe3+ ion in the iron (III) chloride reacts with phenol groups to form a purple complex.
Salicylic acid contains a phenol group, but aspirin does not. Therefore, if you add iron (III) chloride to an aspirin sample and you see a purple colour, it means that there is still some salicylic acid present and the sample is impure. Calculating % yield of my aspirin The chemical formula for the reaction is.. C7H6O3 (salicylic acid) + C4H6O3 (Acetic anhydride) ? C9H8O4 (aspirin) + C2H4O2 (acetic acid) The ratio of salicylic acid to aspirin is the same as acetic anhydride to aspirin: 1 to 1. This means one mole of salicylic acid (or acetic anhydride) makes one mole of aspirin.
Since there is less moles of salicylic acid than acetic anhydride, salicylic acid is the limiting reactant, so you use it to calculate the amount of aspirin. Since the mole ratio is 1:1, 0. 015 mol salicylic acid (n=m/Mr, n=2/138=0. 0145mol) makes . 015 mol aspirin. You then use the Molar mass of aspirin times the number of moles of aspirin (180×0. 015=2. 7g) to calculate the theoretical yield of aspirin which is 2. 7g. Then to calculate the % yield of aspirin you use the equation.. Therefore the % yield for my reaction = (2. 452/2. 70) x 100 = 90. 78% Recrystallization.
Final purification of my aspirin was accomplished by the process of recrystallization. Recrystallization, is a procedure for purifying an impure compound in a solvent. I dissolved my impure aspirin in the solvent ethyl ethanoate to create a highly concentrated solution, which I went on to heat to a high temperature. I used the solvent ethyl ethanoate as aspirin it’s impurities as polar solvents dissolve polar reactants and all of my reactants were polar. According to the principle that recrystallization is based on, the amount of solute that can be dissolved in a solvent increases as the temperature increases.
This is because for the aspirin to dissolve the forces between the particles in it must be broken. This is an endothermic process called dissociation. In addition, some of the intermolecular forces between the particles in the liquid must also be broken also in order to enable bonds to form between the aspirin and the solvent. This process is also endothermic, hence why a high temperature is needed. The particles of the solid and the particles in a liquid become attracted. This again is an exothermic process called solvation. Therefore heating my concentrated solution allowed the aspirin and the impurities to dissolve in the solution.
As the solution cools the solubility of the impurities in the solution and the aspirin being purified decrease. The aspirin should crystallize before the impurities as there was more aspirin in the solution than there were impurities. I then used vacuum filtration to separate the more pure crystals of aspirin. As the impurities were still in solution they passed through the filter paper leaving only the aspirin crystals to collect on the filter paper. Melting point A pure compound will melt over a relatively narrow temperature range, as impurities both lower the temperature and widening the range over which a compound melts.
As a general rule, a 1% impurity will lower the melting point of a sample by 0. 5 °C1. The melting point of Aspirin is 138–140 ° C2. When aspirin is heated, the heat energy that’s added to the substance is translated into kinetic energy. The more kinetic energy a compound has the more the particles it is made up of are able to partially overcome the intermolecular attractive forces which keep the particles held rigidly in place in the highly-ordered crystal structure. The melting point of a substance is the temperature range over which the first crystal of the solid starts to melt and the last crystal completes its melting.
A melting point range is very narrow for pure solids (usually just 1 – 2°C), characteristic of the particular compound. The presence of even a small amount of impurity will lower a compound’s melting point by a few degrees and broaden the melting point temperature range because the impurity causes defects in the crystalline lattice, it is easier to overcome the intermolecular interactions between the molecules. In a very pure organic crystal, all of the molecules are the same, so they pack together in a perfect, very orderly way.
In this arrangement, the attractive forces between the molecules are maximized and uniform. However an impure compound contains more than one organic compound. These different molecules would not fit together properly to make an orderly arrangement. This additional molecule creates a defect in the structure of the crystal, so the structure is weak, and is more easily overcome by an input of energy. Less heat is needed to melt this mixture than is required to melt the pure structure. Therefore, impure aspirin melts at a lower temperature than pure aspirin with no impurities present.
Impure aspirin also melts over a broader temperature range, due to different regions within the crystal structure that contain different amounts of the impurity, and therefore melt at different periods of time. TLC TLC is used to separate substances into their components. TLC involves a stationary phase and a mobile phase. The thin film of silica gel is used as the stationary phase. The mobile phase is designed to be more non-polar than the stationary phase and flows through the stationary phase to carry the components of the test substance up the plate. TLC plates can be one of two kinds, Silica gel or alumina.
I used silica gel TLC plates. Silica gel is a form of silicon dioxide. The silicon atoms that make up the plate are joined via oxygen atoms in a giant covalent structure. However, at the surface of the plate, the silicon atoms are attached to -OH groups. Meaning you have Si-O-H bonds instead of Si-O-Si bonds. Therefore the surface of the silica gel plate is very polar and, because of the -OH groups, can form hydrogen bonds with suitable compounds as van der Waals dispersion forces and dipole-dipole attractions. Different components travel at different rates.
As the solvent begins to soak up the plate, it starts by dissolving the compounds in the spots of the samples you have put on the base line. The compounds present then get carried up the chromatography plate as the solvent continues to move upwards. The compound will get carried up the plate at a faster rate if the compound is very soluble in the solvent, meaning that there is a lot of attraction between the molecules of the compound and the molecules of the solvent.
The compounds will travel slower up the plate if the compound sticks to the silica gel very strongly, i.e. if a compound forms strong bonds with the Si-O-H bonds of the silica gel then they will more slower up the plate. It is very unlikely that all your samples will bond to exactly the same extent, and be soluble in the solvent to exactly the same extent. Also it is good to note that it isn’t just the attraction of the compound to the silica gel that matters. Attractions between the compound and the solvent are also important – it effects how easily the compound is pulled back into solution away from the surface of the plate.
As the solvent slowly travels up the plate, the different components of the dye mixture travel at different rates and the mixture is separated into different coloured spots. If the other tested compounds leave exactly the same pattern on the chromatography plate as the known pure compound we can conclude that they are the same. However, if extra spots are observed as well as the characteristic pattern of the known compound, then impurities are likely to be present in the sample. The reason for not touching the powder side of the TLC plate is because oils from your skin or residue on gloves can contaminate the plate and obscure your results.
The reason you put a lid on the tank is to make sure that the atmosphere in the beaker is saturated with solvent vapour. Saturating the atmosphere in the beaker with vapour stops the solvent from evaporating as it rises up the plate. Fe3+ test Salicylic acid forms an intense purple coloured complex with Fe(III) ions as a consequence of the phenol group in its structure. Pure aspirin gives no colour because the phenol group has been converted into an ester. Iron is a transition metal. All transition metals have at least ion that is coloured. Iron has several coloured ions but I am using the Fe3+ that is purple in colour.
Crystal field theory describes the breaking of d or f orbitals, due to a static electric field produced by anion neighbors (in my case the negative oxygen in the phenolic group of the aspirin). The transition metal complex forms because of the attraction between the positively charged iron cation and negative charge on the non-bonding electrons of the phenolic group on the salicylic acid. As the salicylic acid approaches the iron ion, the electrons from the –OH group will be closer to some of the d-orbitals and farther away from others causing a loss of degeneracy.
The electrons in the d-orbitals and those in the salicylic acid repel each other due to repulsion between like charges. Therefore resulting in the d-electrons closer to the phenolic group having a higher energy than those further away, which results in the d-orbitals splitting in energy. Please see the hydrolysis of aspirin background theory below for more information on this. Hydrolysis of aspirin When aspirin undergoes hydrolysis, the products are salicylic acid and acetic acid. The hydrolysis reaction causes the dissociation of water and results in the replacement of an organic functional group with a hydroxide.
When an O-C bond is broken the –H and –OH of the dissociated water adds to the two molecules where the bond was broken. The phenol group on the salicylic acid produced by hydrolysis reacts with the Fe3+ ions to form a purple complex as shown on the right. Aspirin does not have a phenol group, instead it has the ester group OOCCH3, which means there is no phenolic group for the iron ions to bond to and cause a purple complex. The intensity of the colour depends on the salicylic acid concentration in a sample as there is no purple colour formed between the aspirin and the iron ions because aspirin does not have.
Therefore the amount of salicylic acid produced can be determined from the intensity of the purple complex it forms with Iron (III) Ions as they are directly proportional. The intensity of the purple colour can be measured by colorimetry. White light is made up of many different colours or wavelengths of light. A coloured sample typically absorbs only one colour or one band of wavelengths from the white light. Therefore only a small difference would be measured between white light before it passes through the coloured aspirin sample compared to after it passes through the coloured aspirin sample.
This is because the one colour absorbed by the sample is only a small portion of the total amount of light passing through the sample. However, by selecting the one colour or band of wavelengths of light to which the aspirin sample is most sensitive (green), we are able to see a large difference between the light before it passes through the sample and after it passes through the sample. The colorimeter passed a green light beam through one of four optical filters to the photo detector where it is measured.
The difference in the amount of coloured light transmitted by the sample is a measurement of the amount of coloured light absorbed by the sample. The choice of the correct wavelength for testing is important. The wavelength that gives the most sensitivity for a test factor is the complementary colour of the test sample. For example The Aspirin-Iron (III) chloride test produces a purple colour proportional to the salicylic acid concentration in the sample (the greater the salicylic acid concentration, the darker the purple colour).
I selected a wavelength in the green region to analyse this sample since the purple solution absorbs mostly green light. Esters are susceptible to catalytic hydrolysis by both aqueous acids (pH 4) and bases (pH9. 2). Making my own sample of Salicylic acid (to then be converted into Aspirin) Method 1. Set up reflux apparatus as shown in diagram.
Remove the pear shaped flask and add 2g of Oil of Wintergreen (Methyl Salicylate) directly to minimise percentage uncertainties and make the results more accurate. 3. sMeasure out 25cm?of Sodium Hydroxide 2moldm?? with a pipette and transfer into the pear shaped flask. 4. Add roughly half a spatula of anti-bumping granules. 5. Attach the pear shaped flask to the condenser.
Making sure the joint is air tight. 6. Boil water in a kettle. While waiting for the kettle to boil arrange the rings of the water bath around the funnel of the pear shaped flask and the water bath so no steam can escape, ensuring the maximum temperature is reached and maintained (You may need to remove the pear shape flask to do this). 7.
Poor the boiling water into the water bath and quickly lower the reflux apparatus into the water bath, making sure the pear shaped flask is sufficiently submerged and there are no gaps between the rings. 8. Leave for 30 minutes. 9. Lift the reflux apparatus out of the water bath after 30 minutes and remove the pear shaped flask. 10. Pour the mixture into a small beaker surrounded by ice and water and add conc. Hydrochloric acid (6M)drop wise until it is just acidic, stirring all the time. Determine the acidity of the mixture using indicator paper (a light to medium orange is the colour you want).
11. Set up a buchner funnel and suction apparatus as shown in the diagram. 12. Filter the product. 13. Wash the product with a little ice cold water and transfer it to a pre-weighed watch glass (so you can minus the weight of the watch glass when calculating percentage yield). 14. Weigh the product on the watch glass. 15. Allow product to dry overnight and weigh the dried product. Converting salicylic acid into aspirin Method 1 1. Set up reflux apparatus as shown in diagram. 2. Weigh out 1g of salicylic acid (2-hydroxybenzoic acid) and put it in the dry pear shaped flask.
3. Measure out 2cm?of ethanoic anhydride and add to flask along with 8 drops of conc. Phosphoric acid. 4. Boil water in a kettle. While waiting for the kettle to boil arrange the rings of the water bath around the funnel of the pear shaped flask and the water bath so no steam can escape, ensuring the maximum temperature is reached and maintained (You may need to remove the pear shape flask to do this). 5. Poor the boiling water into the water bath and quickly lower the reflux apparatus into the water bath, making sure the pear shaped flask is sufficiently submerged and there are no gaps between the rings.
6. Leave until all the solid has dissolved and then continue to warm for a further 5 minutes. 7. Lift the apparatus out of the water bath and remove the pear shaped flask. 8. Add 5cm? of cold deionised water to the solution and stand the pear shaped flask in a beaker of water and ice until the precipitation appears to be complete. You may have to stir vigorously with a glass rod to start the precipitation process. 9. Set up a Buchner funnel and suction apparatus while you wait for the precipitation to complete. 10. When precipitation complete filter the product. 11.
Wash the product with a little cold water and transfer it to a pre-weighed watch glass (so you can minus the weight of the watch glass when calculating percentage yield). 12. Weigh the product on the watch glass. 13. Allow product to dry overnight and weigh the dried product. Method 2 1. Measure out 2g of salicylic acid (2-hydroxybenzoic acid) and 4cm? ethanoic anhydride and place in a 100cm? conical flask. 2. Add 5 drops of conc. sulphuric (VI) acid. 3. Agitate for 10 minutes. 4. Set up water bath and fill with ice and water. 5. Add 4cm? of cold glacial ethanoic acid and place the conical flask in the water bath.
6. Set up Hirsch funnel and suction apparatus (same as Buchner funnel and suction apparatus diagram above but with a Hirsch funnel). 7. Filter the product. 8. Wash the product with a little ice cold water and transfer it to a pre-weighed watch glass (so you can minus the weight of the watch glass when calculating percentage yield). 9. Weigh the product on the watch glass. 10. Allow product to dry overnight and weigh the dried product. Recrystallization of my aspirin Method 1. Weigh out 2cm? of my aspirin made earlier in a test tube. 2. Set up a water bath with boiling water from a kettle in.
3. Add 10cm? ethyl ethanoate and place test tube in water bath. 4. While test tube in water bath set up Buchner funnel and suction apparatus. 5. Filter the solution while it is hot. 6. Allow the solution to cool slowly and watch for crystals to form. If you do not see any crystals forming stir vigorously for a few minutes. 7. Once cool and crystals have formed filter the solution again. 8. Wash the product with a small amount of cold solvent and transfer it to a pre-weighed watch glass (so you can minus the weight of the watch glass when calculating percentage yield). 9.
Weigh the product on the watch glass. 10. Allow product to dry overnight and weigh the dried product. Fe3+ test Method 1. Get 3 test tubes and label them recrystallized aspirin, my aspirin and commercial aspirin. 2. Put 1mL of deionized water into each test tube. 3. Add one crystal or a “pinch” of each sample into its test tube. 4. Add one drop of iron (III) chloride solution to each test tube. 5. Record your observations. Measuring melting points of my aspirin, recrystallized aspirin and commercial aspirin Method 1. Set up a Bunsen burner on a heat proof matt. 2. Take three melting point tubes.
Touch one end of each melting point tube to the outside edge of a roaring flame as shown on the diagram. 3. Fill each of the three melting point tubes to the depth of roughly 8mm with each of your three samples (your aspirin, commercial aspirin and recrystallized aspirin). Make sure each sample is dry by checking that it has had at least been left one night to dry. Label each tube by laying it on labelled piece of paper once filled. To fill the melting point tubes it is easiest to dip the open end into the sample and gently flick the sample in the tube to dislodge it so it falls down to the sealed end of the tube.
4. Set up melting point apparatus. 5. Place the melting point tubes in the heating block. Make a note of which sample is in each position. 6. Turn the power switch to on and also the boost switch. We know the melting point range of aspirin is 130-140°C so we can heat the samples to around 100°C. Turn the temperature dial to 10 so it heats up fast. Keep an eye on the temperature so you do not miss the melting point. 7. While the temperature is rising adjust the watch glass so it is in focus, making sure you can see each sample clearly. 8.
When the thermometer reads 100°C turn of the boost switch and allow the temperature to slowly rise to 130°C while watching the three samples through the watch glass. 9. Note the melting point range for each sample. Remember the range is the temperature at which the first liquid droplet forms and the temperature when the last solid particle appears to melt. General TLC method 1. Set up Bunsen burner and stretch out capillary tubes so they have a smaller diameter and give smaller spots. 2. Place a small amount of pure aspirin, recrystallized aspirin and commercial aspirin three separate test tubes and label each.
3. Add 1cm? of ethanol to each test tube to dissolve the samples. You may need to swirl the test tubes to make all the sample dissolve. 4. Lay your TLC plate face up and draw on your base line with a pencil and mark three X’s for where you will be spotting your samples. 5. Use the capillary tubes you stretched out earlier to spot each of your three samples on the TLC plate, draw a diagram on paper so you know which spot is which sample. Allow the spots to dry then repeat a few times. 6. While the spots are drying make your tank by adding a sensible amount of ethyl ethanoate to a 100cm? beaker.
Make sure the ethyl ethanoate will not go above your base line. Place cling film over the beaker. 7. Open the lid of the tank (cling film) and place in the now dry TLC plate and allow the solvent to rise to within a few mm of the top of the plate. 8. While the solvent is rising make your iodine tank by adding a few iodine crystals into a 100cm? beaker and sealing it with cling film. 9. When the solvent has risen mark the solvent front with a pencil and place the TLC plate into the iodine tank. 10. When spots have appeared remove the TLC plate and work out the Rf values if necessary or simply compare the different spots.
Remember the plate will fade. Hydrolysis of aspirin Buffered solutions Method 1. Crush one buffered tablet using a pestle and mortar, the finer it is the faster it will dissolve. 2. Measure out 100cm3 of deionised water using a measuring cylinder. 3. Add the deionised water and crushed buffer tablet to a 150mL beaker and stir until all the tablet has dissolved. 4. Repeat this step for the pH’s 4,7 and 9. 5 by using individual pH tablets each time that correspond with the desired pH. 5. Place each pH solution in a sealed bottle and label it. 0. 0025 mol dm-3 Iron (III) nitrate solution.
Method 1. Weigh out 10g of iron (III) nitrate-9-water into a 250mL beaker. 2. Measure out 50cm3 of deionised water and stir until all the solid has dissolved. 3. Measure out 1cm3 of the solution and transfer to a 7mL bottle. 4. Make the solution up to 5cm3 by adding 4cm3 of deionised water. 5. Repeat this step about 10 times so you have 10 7mL bottles of 0. 0025mol dm-3 Iron (III) nitrate. Hydrolysis of aspirin Method 1. Take your 100cm3 pH 7 solution you made earlier and heat it to 70°C using Bunsen burner, heat proof mat, tripod and gauze apparatus. 2.
While it is heating up set up your colorimeter by plugging it in at your workstation and selecting the digits option and pressing the F4 button to create 4 digit readings. 3. Fill a cuvette with deionised water. 4. Place the curvet in the sample chamber and press the green button to calibrate the colorimeter. When the light on the green button is green and all the absorbance values on the screen are 0 then you have calibrated the colorimeter. 5. Now your colorimeter has been set up go back to your pH 7 solution. When it is at 70°(you may need to place the beaker in a larger beaker of cold water if it has gone over 70°C) add 0.
10g of powdered aspirin and stir until it is all dissolved. 6. While it is dissolving pipette your 5cm3 iron (III) nitrate solution that you made earlier into a cuvette. 7. When all of the aspirin is dissolved in the pH solution measure out 1cm3 of the reaction mixture and pipette it into the same cuvette. 8. Remove the deionised water filled cuvette from the sample chamber in the colorimeter and replace it with the pH 7 aspirin cuvette you just made up. 9. Take a reading from the green absorbance filter immediately. This reading represents time 0 but it is likely some decomposition will have taken place.
Take readings every 30 seconds and record them in a table.
References Synthesis of salicylic acid – hydrolysis http://www. rsc. org/learn-chemistry/content/filerepository/CMP/00/000/045/Aspirin. pdf http://www. chemguide. co. uk/organicprops/esters/hydrolysis. html Synthesising aspirin – esterification http://homepages. ius. edu/DSPURLOC/c122/asp. htm https://eee. uci. edu/programs/gchem/D03MANAspirinSynAnalysis. pdf http://wwwchem. csustan. edu/consumer/aspirincons/aspirincons. htm http://www2. vernier. com/sample_labs/CHEM-A-22-COMP-aspirin. pdf http://homepages. ius.edu/DSPURLOC/c122/asp. htm http://facweb. northseattle. edu/amohamed/CHEM%20102N/Labs/Lab%207.
Esterification. pdf http://chemwiki. ucdavis. edu/Physical_Chemistry/Physical_Properties_of_Matter/Solutions/Case_Studies/RECRYSTALLIZATION Synthesis of aspirin – acetic anhydride + water reaction and equation http://www. laney. edu/wp/cheli-fossum/files/2012/01/8-Synthesis-of-Aspirin. pdf# Recrystallization http://www. laney. edu/wp/cheli-fossum/files/2012/01/8-Synthesis-of-Aspirin. pdf#
Melting points http://www. chm. uri. edu/mmcgregor/chm228/use_of_melting_point_apparatus.pdf https://dlm. chm. bris. ac. uk/demodlm/mp. htm TLC http://chemistry. oregonstate. edu/courses/ch361-464/ch463/TLC_for_Ketones. htm http://www. chemguide. co. uk/analysis/chromatography/thinlayer. html.
Fe3+ test http://en. wikipedia. org/wiki/Crystal_field_theory#Explaining_the_colors_of_transition_metal_complexes Hydrolysis of aspirin – colorimetry http://www2. volstate. edu/chem/1110/Purity_of_Aspirin. htm http://www. omega. com/manuals/manualpdf/M3780. pdf http://infohost. nmt. edu/~jaltig/AspirinSupp. pdf General information on aspirin tablets.
http://faculty. mansfield. edu/akiessli/Organic%20Chem%201/Lab%20handouts/Recrystallization%20of%20Aspirin. pdf Results Synthesis of salicylic acid % yield of salicylic acid = 90. 78% Iron (III) test positive Measuring melting points My aspirin Recrystallized aspirin Commercial aspirin 1 2 3 Mean 1 2 3 Mean 1 2 3 Mean Start m. p. /°C 129 130 130 129. 7 120 121 121 120. 7 127 127 127 127. 0 End m. p. /°C 136 134 135 135. 0 127 130 127 128. 0 134 134 135 134. 3 Hydrolysis of aspirin Results for pH 7 Trial 1 Time/ min Green absorbance reading/ 565nm.