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VI. Results and Discussion In this exercise, the goal was to produce acetylsalicylic acid through the organic synthesis from the reaction of salicylic acid to acetic anhydride, the starting materials. Instead of using acetic acid, acetic anhydride was used as solvent since the anhydride reacting with water to form acetic acid tends to drive the reaction to the right. It results from the elimination of a molecule of water from two molecules of acetic acid (see Fig. 11.2). Figure 11.3 below shows the balanced chemical reaction of the synthesis of acetylsalicylic acid.

Figure 11.2. Structure of Acetic Anhydride Figure 11.3. The balanced chemical reaction of the formation of aspirin. Because the reaction is slow in pure acetic anhydride, the catalyst, commonly strong acids like phosphoric acid was used for the reaction. According to Le Chatelier’s principle, the presence of excess acetic anhydride forces the equilibrium towards the desired product, which in this case is the aspirin. In addition to this, the catalysts were also used to ensure that side reactions, which may cause the percentage yield to increase, will be avoided. The reaction behind the synthesis is nucleophilic acyl substitution. According to McMurry (2000), nucleophilic acyl substitution happens when the initially formed intermediate expels one of the substitutes originally bonded to the carbonyl carbon leading to the formation of a new carbonyl compound. In this experiment, the specific nucleophilic acyl substitution is esterification. It occurs when a carboxylic acid in salicylic acid and an alcohol combine in a reaction to produce an ester.

Figure 11.4. Mechanisms on the formation of aspirin

Phosphoric acid protonates the carbonyl oxygen atom (C=O) of the anhydride to make it more prone to nucleophilic attacks. It gives the anhydride a positive charge thus, making it more susceptible to nucleophilic attacks. The nucleophilic hydroxyl group of salicylic acid attacks the electron deficient acetic anhydride resulting to a tetrahedral intermediate. The hydroxyl group (-OH) attached to the electrophilic carbon removed the hydrogen as proton thus donating the electron to form a double bond (C=O).The loss of the proton regenerates the phosphoric acid and thus, producing acetylsalicylic acid. To enhance the synthesis reaction, addition of heat and water after heating were done. The synthesis of reaction is favored by heating the mixture because it speeds up the dissolution process of salicylic acid and increases its solubility as well. Because this specific reaction is an endothermic process, addition of heat would favor a forward reaction resulting to the formation of products. Aside from that, nucleophile was completely facilitated by the addition of water after heating. Water was used in order to provide medium for further nucleophilic substitution. The theoretical yield obtained is 1.30 grams after knowing that salicylic acid is the limiting reagent. The actual yield obtained is 0.78 grams thus, the % yield is 60% . This results are relatively low because of possible sources of error such as loss of product in the filter paper because of prolonged air dying, decomposition to acetic acid in solution so there wasn 't a complete conversion of reagents and insufficient heating. Upon obtaining the crude aspirin, recrystallization was done. This is performed to remove the traces of impurities. After cooling to room temperature and immersing on an ice cold water bath, suction filtration method was done to separate the filtrate from the residue which contains the recrystallized products. Suction filtration is the most practical technique to use when fast filtration of mixture is desired. It employs vacuum which can aid in the passage of filtrate through the filter paper (Basic Organic Chemistry: Laboratory Manual, 2012). In addition, since aspirin is an ester, it should not be recrystallized from hot water because it will allow the crude sample to be hydrolyzed and yield undesirable products. The % recovery obtained is 59%. After performing the synthesis of aspirin from salicylic acid, the verification of the identity and purity of the product through melting point determination was also performed. In differentiating salicylic acid from the synthesized product, FeCl3 test and KMnO4 test were conducted. For the FeCl3 test, the positive will give a change from yellow-brown solution to a violet colored complex. For this test, the result is positive because of the presence of phenol in it. The oxygen atoms of the acid group –COOH, and of the -OH group on the salicylic acid together can form a complex with Fe(H2O)6 +3. The test result on aspirin is negative because iron complex cannot be formed due to the absence of a hydroxyl group attached to benzene. For the KMnO4 test, a positive result was obtained from the synthesized aspirin as seen by the disappearance of violet color and formation of the brown precipitate. Theoretically, salicylic acid would give a positive result because of the presence of hydroxyl group (-OH). Recalling the reactions in alcohols, KMnO4 was used to detect the presence of primary and secondary alcohols. Since the phenol group is absent in the synthesized aspirin, a negative result should be obtained. Another test was conducted to differentiate the synthesized aspirin from the commercially-available aspirin. For the iodine test, the positive result is the solution turning blue or violet because of the presence of starch. In this particular experiment, it is the commercially-available aspirin that has a color change which is the solution exhibited a faint violet color. Theoretically, this holds true because pharmaceutical companies add starch to tablet medicines to give them its characteristic shapes. Lastly, the melting point determination was conducted. One way of identifying a substance is through its melting point. The range of the melting point can give one an idea on the purity of the sample. The theoretical melting point range of aspirin is 128-137°C. In this experiment, the obtained melting point range of of the crude aspirin is 106- 112°C and for the recrystallized aspirin, it is 114-118°C. The result of melting point determination means that the samples have impurities in it. Narrow difference in the melting point range of the sample and the theoretical melting point range means that the substance is pure because of the uniform forces present in the molecules. When the range is wide, it means that the sample contains impurities. In this exercise, melting point determination is done. And the results show that it has a wide difference from the theoretical melting point range of aspirin. Thus it can be inferred from the results that the sample is not pure. VII. Summary and Conclusion Synthesis of organic compounds involves guidelines and steps that should be followed. The first one is the establishment of the starting materials which is in this are the salicylic acid and acetic anhydride together with the phosphoric acid that served as the catalyst for the reaction. The synthesis of aspirin involved the acid-catalyzed nucleophilic acyl substitution. The specific nucleophilic acyl substitution for this experiment is esterification. It happens when a carboxylic acid from the salicylic acid and an alcohol combine in a reaction therefore producing an ester.
In this experiment, phosphoric acid was used as a catalyst to hasten the reaction between the salicylic acid and acetic anhydride. Heat and addition of water was also done for efficient production of the desired product. The percent yield obtained for this experiment is 60%. Low %yield can be caused by insufficient heating and that the product was lost in the filter paper because of prolonged air drying.
The next step that was done was the recrystallization of the crude sample to obtain a more purified organic compound. Recrystallization was done by suction filtration. The last step for this experiment is the verification and differentiation of the samples. Through the KMnO4 test, it was verified that the synthesized product was indeed to be acetylsalicylic acid or most commonly known as aspirin. Another test that was conducted is the iodine test wherein the commercially-available aspirin gave the positive result due to the presence of starch in it. Unfortunately, FeCl3, did not give the theoretical results.

VIII. References Aspirin timeline. (2013). Retrieved on October 5, 2013 from http://www.telegraph.co.uk/health/healthnews/8185164/Aspirin-timeline.html

History of Aspirin. (2013). Retrieved on October 5, 2013 from http://en.wikipedia.org/wiki/History_of_aspirin

Institute of Chemistry. (2012). Basic Organic Chemistry Laboratory Manual. University of the Philippines Los Baños College Laguna. 72-75. Material safety data sheet. (2013). Retrieved on October 6, 2013 from http://www.inchem.com.ph/productpages/fecl3_msds.pdf

McMurry, John. (2000). Organic Chemistry. 7th Edition USA.

MSDS for potassium permanganate. (2013). Retrieved on October 6, 2013 from http://www.sciencelab.com/msdsId=9927406

Phosphoric acid. (2013) retrieved on October 6, 2013 from http://www.ccohs.ca/oshanswers/chemicals/chem_profiles/phosphoric.html

Rodriguez, E.B. (1997). Basic Principles of Organic Chemistry. UP Open University: Diliman Quezon City . 295 – 336.

IX. Remarks and Recommendation The synthesis of organic compounds is indeed very helpful in chemistry and through this process, one can have a glimpse on how chemical processes works in real life.
The researcher recommends the use of other tests in order to obtain more accurate results in the differentiation of the synthesized aspirin from the commercially-available aspirin.

I. RESULTS AND DISCUSSIONS

On the first part of the exercise, acetylsalicylic was produced by reacting salicylic acid and acetic anhydride with the presence of a catalyst, phosphoric acid. The product was a white pile of needle like crystals with minimal colorless liquid. The mechanism of the reaction is shown below:

Figure 11.3 Mechanism for the reaction involved in Synthesis of Aspirin

Heating was employed in mixture of salicylic acid and acetic anhydride to effectively dissolve the salicylic acid and aid in the reaction of salicylic acid and the acetic anhydride. The addition of water enabled the decomposition of acetic anhydride after the formation of aspirin. This addition of water was not done at the start of the experiment since this will initiate a reaction with the anhydride, forming acetic acid thus, the reaction between the acetic anhydride and salicylic acid will not push through.

On the other hand, phosphoric acid donates an H+to the reaction complex to speed up the reaction. Note that H+ will not be consumed in the reaction since it is a catalyst. When acetyl chloride is used instead of acetic anhydride, Phosphoric acid will not be needed anymore since the reaction will push through by Friedel-Crafts Acylation.

After letting the reaction upon addition of water finish, 20 mL of ice-cold distilled water was added to the mixture to start lowering the temperature and initiate crystallization of the product. The resulting mixture was then filtered through suction filtration. Suction filtration was used since the solid particles of the mixture would not let the liquid through if simple filtration set up was used, instead.

The collected crude product (residue) was then washed with distilled cold water several times and then air dried.

The crude aspirin still contains a lot of impurities thus, it was subjected to recrystallization. It was first dissolved through the use of ethanol. And then cold distilled water was added. Cold water is used so that the crude product will recrystallize. Esters have relatively low melting points thus the use of hot water would not aid the recrystallization of the product. The mixture was subjected to cool water bath. Note that it is cool water bath that is used instead of cold water bath. Gradual decrease in temperature is essential since this will enable the product to slowly recrystallize and will not cause impurities to be trapped in the lattice of aspirin. The mixture was then again filtrated through suction filtration. Product was air dried.

The collected recrystallized aspirin was then subjected to various chemical tests to differentiate it from the starting material and to characterize it. Aspirin and salicylic acid was tested with potassium permanganate and ferric chloride. On the potassium permanganate test, aspirin caused the purple color of the solution to lighten slightly while salicylic acid caused the purple color to fade and then formed orange-brown precipitate which is the positive result for phenols. The same results were obtained in the ferric chloride test. Aspirin only showed slight change of color to light blue complex while salicylic acid produced a dark blue violet complex. Both results of aspirin in the two tests showed that there is a presence of a compound containing the phenol functionality, though minimal, in the product. This may be due to insufficient air drying; minimal amounts of the starting material may have been trapped in the compound since it was not fully dried when subjected to the chemical tests.

In the iodine solution test, aspirin just settled at the bottom and some dissolved while the commercially available aspirin formed black precipitate which indicates the presence of starch.

II. SUMMARY AND CONCLUSION

Aspirin was synthesized by reacting salicylic acid and acetic anhydride with the presence of a catalyst. The product collected was a milky-whitepile of needle like crystals. The collected product was recrystallized to remove impurities such as unreacted acetic anhydride and salicylic acid. The recrystallized product was subjected to various chemical tests to differentiate it from the starting materials. FeCl3 and KMnO4 tests indicated the presence of phenol functionality in the compound. This may be due to insufficient washing and air drying. In the iodine test, the result indicated that in the commercially available aspirin, there is the presence of starch.

Though the compound reacted with reagents for testing phenols, it was assumed that the identity of the compound was acetylsalicylic acid since the reaction was not as strong as when the reagents were reacted with salicylic acid.

III. REFERENCES
Baum, Stuart J. and Hill, John W. 1993.Introduction to Organic and Biological Chemistry. New York : Macmillan Publishing Company, 1993. p. 121.
Scott, Ronald M. 1980.Intoduction to Organic and Biological Chemistry. San Francisco : Harper & Row, Publishers, Inc., 1980. p. 189.

http://oehha.ca.gov/public_info/pdf/Iodine%20Fact%20Sheet%20Meth%20Labs%2010 '03 '.pdf

IV. RECOMMENDATIONS
It was noticed that in the laboratory experiment, time is not enough. It would be better if it will be given two meetings to finish the experiment to give way for sufficient air drying of sample since this is one of the most common sources of error.

VI. RESULTS AND DISCUSSION Aspirin can be made by using a process called esterification. Esterification occurs when a carboxylic acid and an alcohol combine in a reaction to produce an ester. This reaction can be used to synthesize aspirin from salicylic acid. In the lab, the carboxylic acid alcohol mixture is heated in the presence of H2PO4, phosphoric acid, which acts as a catalyst. During the reaction process, a molecule of water splits off and the remaining carboxylic acid and alcohol fragments become attached producing an ester.

Fig. 11.5 General Reaction of Esterification

Fig. 11.6 Esterification of Aspirin Usng Acetic Acid The first part of the experiment was the preparation of Acetylsalicylic Acid (Aspirin). A white, milky mixture was obtained when salicylic acid, acetic anhydride and phosphoric acid (a catalyst) were mixed. The mechanism of the reaction is:

Fig. 11.7 The synthesis of Aspirin
For further explanation, an acetic anhydride instead of acetic acid was used because acetic anhydride has a faster reaction time. The catalyst concentrated H2PO4 is added to speed up the reaction. The flask containing this solution is then heated in a boiling water bath for about 15 minutes. This process is called esterifcation.

Fig. 11.8 Esterification of Aspirin In Our Lab Using Acetic Anhydride
The flask is then removed and allowed to cool. Ice cold distilled water is slowly added to the flask to decompose any unreacted acetic annhydride.

Fig. 11.9 Decomposition of Unreacted Acetic Anhydride This shows that the oxygen in salicylic acid attacks one of the carbons in acetic anhydride. Also, the mechanism shows how acetic acid was separated from the acetylsalicylic acid.

In the first part of the experiment, heating of the mixture was done and a clear yellow liquid was obtained (Table 2). Heating was employed so that salicylic acid would melt and react with acetic anhydride. On the other hand, water was added after heating (not at the start of the experiment). This is to prevent the reaction of acetic anhydride with water at the start of the experiment, if this had happened, no aspirin could have formed. In this manner, acetic anhydride was decomposed after the formation of aspirin.
After the adding 40mL ice-cold water, cooling to room temperature and placing in an ice bath, the liquid became whitish / cloudy with white precipitates. This addition of cold water is very important in purification and isolation of the crystals from the liquid since aspirin is insoluble in cold water. Ice cold distilled water is added to the flask again. The flask is then chilled in an ice-water bath for about 10 minutes until crystallization of the aspirin is complete. The aspirin crystals are collected on a Buchner funnel and washed with additional ice cold distilled water. ice cold distilled water is used instead of room temperature distilled water because aspirin is insoluble in cold water and would not be dissolving any of the aspirin product. The acetic acid and H2PO4 are water soluble, in any temperature water, and can be removed by washing the aspirin with the chilled water. Salicylic acid is only slightly soluble in water and any unreacted salicylic acid cannot be removed completely in the washing process. Once the aspirin crystals are purified they are allowed to dry. They are then weighed and tested for purity.
Purification is needed to eliminate any salicylic acid and acetic anhydride that did not react, as well as the acetic acid product and phosphoric acid. Isolation was done through suction filtration, white, sugar-like crystals were obtained.

The crude/impure product was then weighed and it weighed 1.28 g . This is quite far from the theoretical yield because it still contains impurities. This data was used to calculate the percent recovery on the latter part of the exercise.

The second part of the experiment was recrystallization. This is the second part of the purification process. Here, 95% ethanol was added dropwise to the crystals until dissolved and after this, distilled water was added dropwise until cloudy/until recrystallization. Ethanol was used to dissolve aspirin along with the impurities such as salicylic acid and others. Cold water, on the other hand, is used to recrystallize only aspirin, thus, leaving all the impurities behind. Since aspirin is an ester, it should not be recrytallized from hot water since esters hydrolyses in hot water. After cooling in an ice bath (which further facilitates recrystallization and purification), the mixture was then suction filtered.

On the other hand, the calculated percent recovery was 98.1369%. The weight of the recovered sample was 2.85g. The calculation for percent yield was shown in Table 11.6. The percent yield was _____%, meaning there was a slight error. Perhaps, the sample was not weighed properly or it was weighed when still wet.

As for the melting point data, the range of the crude sample was ______˚C and the range of the purified sample was ______˚C. Comparing the results to the literature value of 135˚C, both the purified and crude had a precise value BUT since the purified sample has a narrower range, it is logically more comparable to the literature value.
The computed % recovery of the recrystallized aspirin from the formula: % recovery = (mass of crude aspirin/mass of recrystallized aspirin) x 100, is 41.19%.

The melting point of the crude and recrystallized aspirin was determined using a MP determination test. Results showed that crude aspirin had a wider MP range of 119-135°C compared to 130-135°C MP range of the recrystallized aspirin. This means that there is still more impurities in the crude than in the recrystallized aspirin.

In differentiating salicylic acid to acetylsalicylic acid, ferric chloride was used because compounds containing phenol group will generate a colored complex such as blue red green or purple. In the laboratory, a purple complex was observed in salicylic acid because it has a phenol group.

In Table 11.7 and 11.8, the differentiation of synthesized acetylsalicylic acid from commercially available aspirin was accounted for. The test used in this part was Iodine test, which is a test for the presence of starch (since iodine can form a black complex with starch).After dissolving synthesized aspirin in 2mL water and 1mL iodine solution, a mixture of red-orange liquid and white precipitates was obtained while when commercially available aspirin was dissolved in 2mL water and 1mL iodine solution, a black precipitate in a dark brown to black solution was formed. This shows that commercially available aspirin contains starch.
In the characterization of Aspirin, an iron III chloride test was done to determine the purity of our aspirin. Iron III chloride combines with the phenol group to form a purple complex. If salicylic acid is present (impurity) the product will turn purple when FeCl3 is added, because salicylic acid is a phenol.
Purple color was observed when FeCl3 added which is a sign of an impure product, therefore, salicylic acid may have been present. Salicylic acid is NOT water soluble and cannot be removed by using cold distilled water. If the solution is heated long enough the reaction will not be complete.
No color change and appearance of yellow color, or a faint purple tinge signifies a pure product, it means no unreacted salicylic acid is present.

Fig. 11.10 LEFT - pure (no phenol present), RIGHT - impure (phenol present) Other tests that were performed were summarized in Table 11.7. Since salicylic acid has a phenol group, it gave a positive result to FeCl3 Test and KMnO4 Test, both of which react with phenol. Acetic anhydride gave a positive result to water solubility test to form acetic acid. The recrystallized aspirin, an ester, did not give any positive result to the tests since esters do not react with FeCl3 Test, KMnO4 Test and Tollens’ Test. Small esters are actually fairly soluble in water but solubility falls with chain length and hydrophobic parts. Since aspirin has a hydrophobic aromatic ring, it did not dissolve in water. Having these results, the recrystallized sample was then identified (or assumed) as acetylsalicylic acid.

The synthesized acetylsalicylic acid was differentiated from commercially available aspirin using iodine solution. Commercially available aspirin have small amount of inert binding material such as starch so it reacted with iodine and form a blue-black colored solution. However, in the synthesized aspirin there was no reaction observed which indicates that the generated product was pure.

The low yield and purity that we obtained in this experiment are not only attributed to the experimental procedure and mechanism, but also to other experimental flaws. The low yields could be due to several factors, such as disturbance during crystallization, incomplete reaction and overheating. After all, crystal formation is highly dependent on critical temperature and the level of disturbance the saturated solution was subjected to. Similar to previous experiments, the usage of filter paper containing paper pulp reduced experimental yields as some product was left behind in the Buchner funnel and on the filter paper after suction filtration. Some were also stuck amongst the fibers of the filter paper that was utilized.
VII. SUMMARY AND CONCLUSION

Aspirin was prepared from the reaction of salicylic acid and acetic anhydride. Phosphoric acid was used as a catalyst. It was heated to have a higher rate of reaction. The mixture was then cooled for the material to undergo crystallization. Upon addition of cold water, acetic acid was formed and thus eliminated. The crystals were separated through suction filtration. The other impurities, such as salicylic acid, from the preparation of aspirin were removed in the process of recrystallization.

The melting point range of the purified and crude samples were compared to the literature value and it showed that the purified sample is logically “near” to the literature value because of its narrow range.

The recrystallized product was differentiated from commercial aspirin through iodine test and it showed that the commercial aspirin contains starch. Other tests such as water solubility test, FeCl3 Test, KMnO4 Test and Tollens’ Test differentiated the starting materials, salicylic acid and acetic anhydride, from aspirin.

Acetylsalicylic acid was differentiated from salicylic acid through ferric chloride test. This test was used to observe the presence of phenol group in a compound. Since salicylic acid has a phenol group, it reacted with ferric chloride, producing a purple complex. The synthesized aspirin did not react because it has no phenol group.

Synthesized aspirin was compared to commercially produced aspirin through the iodine solution test. Commercially produced aspirin is not pure acetylsalicylic acid, it contains binders such as starch will react to iodine solution and produce a blue- black solution.

VIII. REFERENCE

Aspirin(Acetylsalicylic acid). Retrieved September 23, 2012 from http://homepage/smc/edu/gallogly_ethan /files/Aspirin%20Synthesis.pdf

Brown, Lemay and Bursten. 2009. 11th Ed. Chemistry: The Central Science. Prentice Hall, 534-537 Division of Organic Chemistry and Natural Products. (2004). Basic Organic Chemistry Laboratory Manual. 7th edition. UPLB : Institute of Chemistry, College of Arts and Sciences.

McMurry, R.S. 2008. 7TH Ed. Organic Chemistry. USA: Prentice Hall, 332-345.

Norman, R. C. and Coxon, J. (1993). Principles of Organic Syntheses. London, UK: Chapman & Hall.

Schmid, G. H. Organic Chemistry. (1991). St. Louis, Missouri: Mosby-Year Book, Inc.

IX. REMARKS AND RECOMMENDATIONS

During the synthesis and recrystallization of acetylsalicylic acid, it is preferable to maintain the temperature at a fixed 90°C. This is because studies carried out have determined that this is the critical temperature for acetylsalicylic acid crystal formation. At this temperature, the formation of acetylsalicylic acid crystals is at its optimum. Maintaining a certain temperature is almost impossible when dealing with an oil bath on a hot plate. Thus, more advanced heating technology could be utilized, such as heating mantles and heat controllers. Even the common water bath can be used. Unlike the hot plate, the temperature of a water bath is not prone to fluctuations, and a water bath provides us with the exact temperature of the water it contains, unlike a hotplate.

References: Aspirin timeline. (2013). Retrieved on October 5, 2013 from http://www.telegraph.co.uk/health/healthnews/8185164/Aspirin-timeline.html History of Aspirin Institute of Chemistry. (2012). Basic Organic Chemistry Laboratory Manual. University of the Philippines Los Baños College Laguna. 72-75. McMurry, John. (2000). Organic Chemistry. 7th Edition USA. MSDS for potassium permanganate. (2013). Retrieved on October 6, 2013 from http://www.sciencelab.com/msdsId=9927406 Phosphoric acid Rodriguez, E.B. (1997). Basic Principles of Organic Chemistry. UP Open University: Diliman Quezon City . 295 – 336. Baum, Stuart J. and Hill, John W. 1993.Introduction to Organic and Biological Chemistry. New York : Macmillan Publishing Company, 1993. p. 121. Scott, Ronald M. 1980.Intoduction to Organic and Biological Chemistry. San Francisco : Harper & Row, Publishers, Inc., 1980. p. 189.

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