In as they do was not given until

    In the seventeenth century, an Irish chemist called Robert Boyle was the first to label acids and bases on their characteristics, but a reason as to why they behave as they do was not given until 200 years later. Only in the late eighteenth century, a Swedish scientist by the name of Svante Arrhenius proposed that acids are hydrogen-containing compounds that dissolve in water to release protons (H + ) into the solution. Svante Arrhenius also dictated that bases are compounds that dissolve in water to discharge hydroxide ions (OH – ) into the solution. With the Arrhenius definition of acids and bases and Boyle’s observations, the term neutralization came to be as a base can cause an acid to be weaker and contrariwise. Since acids release hydrogen ions and bases release hydroxide ions, when mixed they form water and salt. Even though this was a remarkable discovery at that time, the Arrhenius definition was lacking the reason as to why some bases do not have a hydroxide ion. This was resolved in 1923 by a Danish scientist known as Johannes Brønsted, and an Englishman called Thomas Lowry who independently redefined the Arrhenius definition. The Bronsted-Lowry definition proposed that any compound, which donates a hydrogen ion is an acid and any compound that accepts a hydrogen ion is a base, which covered the reason as to why a does not need to have a hydroxide ion. Through the Brønsted-Lowry definition, acids, and bases are coherent with the concentration of hydrogen ions. Acids can increase the concentration of hydrogen ions since they donate them and bases are vice versa since they accept them.  Acids and bases are classified by the degree of which they ionize in water. Strong solutions ionize fully in aqueous solutions (many ions are formed when dissolved) whereas weak solutions only slightly ionize in aqueous solutions. When an injury occurs the body produces prostaglandins which comes from the parent molecule arachidonic acid. The body then converts that acid to the first prostaglandin PGG2 and from that PGH2. With PGH2 the body can then use it for several other prostaglandins which are responsible for pain, inflammation, blood pressure regulation and blood clotting. Aspirin helps with the lessening of these bodily processes by inhibiting an enzyme that is key to the production of prostaglandins called cyclooxygenase. The active ingredient in aspirin, salicylic acid was discovered as salicin when it was extracted from willow trees in the middle ages. However, salicylic acid irritates the stomach so a german scientist by the name of Felix Hoffmann solved the problem by replacing the hydrogen on the hydroxide groups with C(O)CH3. This created a compound called acetylsalicylic acid which is now what is used in aspirin.     The stomach naturally produces acid in order to aid with the digestion of food and kill bacteria. Since the acid is corrosive and can harm the body, the stomach develops a mucous which protects the stomach lining from being eroded. When too much acid is produced, the stomach lining can begin to breakdown allowing for the actual stomach to get damaged and the forming of ulcers to begin. In some cases, the sphincter (the top of the stomach which keeps it sealed) may not function correctly, permitting the stomach acid to flow into the esophagus causing that burning sensation. The purpose of antacids are to neutralize the excess acid produced in the stomach. Antacids are composed of aluminum hydroxide, magnesium carbonate and magnesium trisilicate, which are all bases and alginates. This results in a neutralization reaction which makes the stomach less acidic and in turn less corrosive. The alginates create a gel which acts as barrier on the sphincter to prevent the acid from coming up. Combined, the antacids aid in pain remediation from ulcers and heartburn. The purpose of this lab was to observe the differences between regular, buffered and enteric aspirin in a simulated stomach environment. In the water solution, the buffered and regular aspirin will quickly dissolve, the enteric will not demonstrate any effect. In the acidic solution, the buffered and regular aspirin will dissolve rapidly, and the enteric aspirin will not dissolve. In the basic solution, all three types of aspirin will dissolve swiftly, with enteric taking the most amount of time. This is so because the purpose of the enteric aspirin is to prevent it from dissolving either in the esophagus or intestines, the buffered capsule will neutralize the acidity of the acetylsalicylic acid.In this experiment, three 250 ml beakers were cleaned and rinsed using soap and water, then filled with 100 ml of water. Using a litmus paper, the initial pH was tested using a litmus paper and universal indicator for each of the beakers, simultaneously, a regular, buffered and ½ of an enteric aspirin using a knife, were added to each of the beakers and a stopwatch for 30 seconds was set. During the 30 seconds, observations were made and the final pH was recorded. The timer was set for another 30 seconds, this was done until a total time of 2 minutes was up. This process was repeated with 100 ml of vinegar and a solution of 1g of baking soda dissolved in 100 ml of water. In Appendix A, table 1, the pH of each type of aspirin increased with the buffered being by the greatest amount and the least being the enteric. In the water solution, the enteric aspirin remained at a pH of 6 while buffered and regular aspirin were reduced with regular being the lowest. The regular aspirin reduced the pH to 6 and the buffered to 7 in the basic solution. In table 2, the regular aspirin dissolved the quickest and fully decomposed while the enteric barely any changes other than a slight discolouration of the solution.The resulting pH after the introduction of the different types of aspirin were as expected. The regular and buffered aspirin reduced the pH of the distilled water due to the acetylsalicylic acid dissociating in water to release more hydronium ions into the solution. The enteric did not have an effect since the capsule is covered to prevent it from dissolving in water as anticipated. The regular and enteric aspirin did not change the pH of the acetic acid since the acetylsalicylic acid could not donate any protons and neither could the acetic acid and the fact that the enteric did not fully break down. The buffered increased the pH due to the neutralization of the buffer in the aspirin, however the strength of the buffer was not high enough to fully neutralize the solution since of the acetylsalicylic acid in the aspirin and the acetic acid in the solution. In the basic solution, the all types of aspirins reduced the pH, with regular being by the biggest difference since there was no buffer to neutralize the reaction and was not enteric coated and therefore decomposed. The fact that the enteric aspirin was expired may have affected the results during the experiment. When aspirin expires it decomposes to form acetic acid, which is weaker than acetylsalicylic acid and can therefore result in a higher pH than expected. Even though the potency of the drug may not decrease significantly after the expiration date, for all purposes of the experiment choosing one that is before the expiry date will result in  more accurate data. Since the enteric aspirin was also cut in half, one side of the capsule would not have the enteric coating. This would cause the pH to decrease in all solutions since with half of the coating gone, the acetylsalicylic acid is exposed and can react with each of the solutions. To extend the lab, determining how much aspirin would it take to cause severe damage to the stomach, already the prescribed amount may cause stomach irritation or bleeding. Aspirin can lower the pH of the stomach but at what point does the stomach lining begins to break down can be tested.

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