The oxygen of (i) uses two electrons to abstract an acidic hydrogen. The single bond breaks and two electrons go to oxygen, forming a carbocation and acetic acid (ii). The carbocation is then attacked by the oxygen of the hydroxyl group of salicylic acid to form (iii). O-H bond breaks and two electrons go to oxygen forming the final product (iv).
1. Why does aspirin dissolve in aqueous sodium hydrogen carbonate whereas the key contaminants do not
By dissolving sodium bicarbonate, the ionic strength of water increases. Sodium ions surround water molecules, preventing the water molecules from interacting with (and thereby dissolving) polar substances. This results in the polar substances (in this case, the impurities) to become less soluble and thus be easily filtered from the desired product. Conversely, the high ionic strength helps dissolve the partially-soluble acetylsalicylic acid in water. It is then possible to recrystallize the sample and expect a relatively higher purity.
1. Explain the chemistry involved in this limit test.
The ferric chloride reacts with phenols and enols, producing a temporary or permanent coloration. Since the salicylic acid itself is a phenol, the ferric chloride test will yield a positive result. The intensity of the coloration is determined by the quantity of the enol or phenol in the sample. The greater the enol or phenol in the sample, the more intense the color. Should the intensity be greater than that of the standard solution, then the sample has a high concentration of salicylic acid which in turn indicates that only a very small amount of the target acetylsalicylic acid has be synthesized.
2. Explain the presence of acetic acid in the standard solution.
Since it is a by-product of the synthesis of acetylsalicylic acid, the presence of trace amounts of acetic acid in the sample is likely. It could also result from the incomplete reaction of the salicylic acid with the anhydride. The anhydride was converted to acetic acid but did not react with the salicylic acid to form the target compound. Acetic acid is then added to the standard solution to control this particular variable.
3. Discuss the results obtained with reference to the methods employed.
A yield of 71.2 was observed for the synthesis of the crude product. Loss was likely from the filtration process that followed after the product crystallized. It could also be due to the incomplete reaction of the anhydride and salicylic acid with phosphoric acid. Furthermore, the latter could explain why the melting point for the crude product was lower than expected.
The observed melting point for the crude product was much lower than the known melting point for acetylsalicylic acid. It is very plausible that this is due to the polymeric by-products mixed with the crude sample. The crude product contained impurities that had relatively low melting points and were of greater quantity than the target compound.
The recrystallized product yield was 23.3, much lower than the yield of the crude product. Since the crude product was subjected to purification process, the unwanted by-products were separated from the target compound thereby lowering the weight of the recrystallized sample. It is also plausible that some of the target product did not dissolve in the aqueous sodium carbonate solution because the ionic strength was not high enough to dissolve salicylic acid a compound only partially-soluble in water. Also, the recrystallization process was only allowed for a brief amount of time so it is likely that some of the product is still in the flask, not yet recystallized.
The recrystallized product was observed to have a melting point approximating that of pure acetylsalicylic acid. The slight deviation from the known melting could be due again to impurities. But since the polymeric by-products have been filtered from the sample, there should be only trace amounts left. The trace amounts could still melt at a lower temperature than the deisred product.
1. Why does aspirin dissolve in aqueous sodium hydrogen carbonate whereas the key contaminants do not
By dissolving sodium bicarbonate, the ionic strength of water increases. Sodium ions surround water molecules, preventing the water molecules from interacting with (and thereby dissolving) polar substances. This results in the polar substances (in this case, the impurities) to become less soluble and thus be easily filtered from the desired product. Conversely, the high ionic strength helps dissolve the partially-soluble acetylsalicylic acid in water. It is then possible to recrystallize the sample and expect a relatively higher purity.
1. Explain the chemistry involved in this limit test.
The ferric chloride reacts with phenols and enols, producing a temporary or permanent coloration. Since the salicylic acid itself is a phenol, the ferric chloride test will yield a positive result. The intensity of the coloration is determined by the quantity of the enol or phenol in the sample. The greater the enol or phenol in the sample, the more intense the color. Should the intensity be greater than that of the standard solution, then the sample has a high concentration of salicylic acid which in turn indicates that only a very small amount of the target acetylsalicylic acid has be synthesized.
2. Explain the presence of acetic acid in the standard solution.
Since it is a by-product of the synthesis of acetylsalicylic acid, the presence of trace amounts of acetic acid in the sample is likely. It could also result from the incomplete reaction of the salicylic acid with the anhydride. The anhydride was converted to acetic acid but did not react with the salicylic acid to form the target compound. Acetic acid is then added to the standard solution to control this particular variable.
3. Discuss the results obtained with reference to the methods employed.
A yield of 71.2 was observed for the synthesis of the crude product. Loss was likely from the filtration process that followed after the product crystallized. It could also be due to the incomplete reaction of the anhydride and salicylic acid with phosphoric acid. Furthermore, the latter could explain why the melting point for the crude product was lower than expected.
The observed melting point for the crude product was much lower than the known melting point for acetylsalicylic acid. It is very plausible that this is due to the polymeric by-products mixed with the crude sample. The crude product contained impurities that had relatively low melting points and were of greater quantity than the target compound.
The recrystallized product yield was 23.3, much lower than the yield of the crude product. Since the crude product was subjected to purification process, the unwanted by-products were separated from the target compound thereby lowering the weight of the recrystallized sample. It is also plausible that some of the target product did not dissolve in the aqueous sodium carbonate solution because the ionic strength was not high enough to dissolve salicylic acid a compound only partially-soluble in water. Also, the recrystallization process was only allowed for a brief amount of time so it is likely that some of the product is still in the flask, not yet recystallized.
The recrystallized product was observed to have a melting point approximating that of pure acetylsalicylic acid. The slight deviation from the known melting could be due again to impurities. But since the polymeric by-products have been filtered from the sample, there should be only trace amounts left. The trace amounts could still melt at a lower temperature than the deisred product.
No comments:
Post a Comment