Added Salicylic Acid To Record The Mass Of Aspirin

Ever wonder about the chemical makeup of tablets that people take for pain relief? Before a tablet can be successfully made, the limiting and excess reactants must be considered. The limiting reactant will affect the amount of the product that can be made. Another reason why the starting reactants must be determined carefully is to make reduce the amount of the reactant in excess so that reactants are not wasted. This experiment uses an Alka-Seltzer tablet. Alka-Seltzer dissolves in water and is an antacid and a pain reliever1. The Alka-Seltzer tablet has many uses such as relief of headaches, ingestion, heart burns, or even upset stomachs2. The active ingredients in an Alka-Seltzer tablet is aspirin, also known as acetyl-salicylic acid (C8H12O4), citric acid (C6H8O7), and sodium bicarbonate (NaHCO3)2. The aspirin in the Alka-Seltzer tablet helps with pain relief. Because of the acid-base chemistry (Brønsted-Lowry), citric acid and sodium bicarbonate produce O2, which makes the tablet fizz when it is dropped in liquid. The Brønsted-Lowry theory shows how the Brønsted-Lowry acid donates a hydrogen ion while the Brønsted-Lowry base accepts the hydrogen ions3. The remaining NaHCO3 that is in excess post reaction with the citric acid is what is used to neutralize stomach acid which helps relief heart burn2. The problem in

In experiment two, the drug Panacetin was separated by a series of chemical reactions into its three components: sucrose, aspirin, and an unknown active ingredient, either acetanilide or phenacetin. The purpose of this lab was to determine what percentages of each component is present in the pain-killer. The initial step was to dissolve Panacetin in dichloromethane. However, sucrose is insoluble in dichloromethane because organic molecules are soluble in organic solvents, and dichloromethane is an inorganic solvent, so only aspirin and the unknown dissolved. By using gravity filtration, sucrose was filtered from the solution and 0.30g of solid was collected.

Both Aspirin and the Unknown are soluble in dichloromethane, due to their non-polar characteristics. To separate the two components, sodium bicarbonate was added (see figure 3). Sodium bicarbonate reacted with aspirin and converted it to a salt, also forming water and carbon dioxide. It was observed that the solution "fizzed" when this reaction took place, showing the release of carbon dioxide. The newly formed salt then traveled to an aqueous layer where it was soluble, while the unknown remained in the dichloromethane layer. The two layers were then separated. To collect an aspirin solid, the combination of the addition of HCl and the process of vacuum filtration helped to break down the salt and form a solid. Then the solid was placed in the Fisher Scientific Biotemp Oven to dry it to a constant mass of 0.091 g, 32.97% of the total composition. The

Not all of the aspirin will be converted to sodium acetylsalicylate, so it will instead remain in the organic layer and be weighed with the unknown. Therefore, the reported weight of aspirin will be too low, while the unknown will be weighed too high.

From the vial labeled “AE,” aspirin (0.533 g) was placed into a 50 mL Erlenmeyer flask with a boiling stick. Toluene (20 mL) was brought to a boil on a hot plate. The boiling toluene (10 mL) was then added to the aspirin until the solid dissolved completely. After allowing the solution to reach room temperature, the solution was placed in an ice bath for 16 minutes. After the crystals

Industrial production of Aspirin is dependent on the company producing the drug; however the general method of production is known as slugging or dry-granulation. In this method, corn starch is mixed with pure water and is then heated and

Separation and Purification of the Components of an Analgesic Tablet. Cora Bruno, Lab Section E. Aspirin, Caffeine and Acetaminophen were separated from four analgesic tablets of Excedrin using extraction techniques. 5% wt/vol NaHCO3, 4M HCL, ethyl acetate and deionized water were used to separate the three active components. MgSO4 was used to dry each extraction. Aspirin was isolated using a hot water bath and weighed to determine the percent theoretical recovery and the actual percent recovery of aspirin. After separation, Aspirin (ASA), Caffeine (CAF), and Acetaminophen (ACE) were purified and identified using Thin Layer Chromatography (TLC). Standards and purified ASA, CAF, and ACE were spotted on the silica gel (stationary phase) of the

Acetic Anhydride and p-Aminophenol were heated in a vial attached to an air condenser to synthesize crude acetaminophen, resulting in 0.097 grams (47.48% yield). The crude acetaminophen was then recrystallized in a solvent of water and methanol over heat resulting in 0.082 grams (39.61% yield) of pure acetaminophen. Melting points of both crude and pure acetaminophen were taken, and found to be 165.9 - 170.9°C and 168.2 - 171.5°C, respectively. The literature melting point of acetaminophen is 169.5 – 171.0°C, indicating that our final product was pure.

The purpose of this lab was to synthesize aspirin, determine the theoretical yield, compare the percent yield to the theoretical yield and test the purity of aspirin by adding Iron (III) chloride to the product.

The goal of this experiment was to synthesize aspirin. In this experiment aspirin, also known as acetylsalicylic acid, was synthesized from salicylic acid and acetic anhydride. In the reaction the hydroxyl group on the benzene ring in salicylic acid reacted with acetic anhydride to form an ester functional group. This method of forming acetylsalicylic acid is an esterification reaction. Since this esterification reaction is not spontaneous, sulfuric acid was used as a catalyst to initiate the reaction. After the reaction was complete some unreacted acetic anhydride and salicylic acid was still be present in the solution as well as some sulfuric acid, aspirin, and acetic acid. Crystallization, which uses the principle of

The only information known about the aspirin tablets used in this lab is they contained ASA, but no further information was known about the pills. For example, we did not know if the tablets had anti-acid properties, which would have affected our data for the amount of titrant needed in order to have a pH change in the sample, as the anti-acid would have counteracted some of the acid added to the sample before titration started. In the lab the only information we were given was that the tablets contained ASA, meaning that all the tests were assuming that the only chemical in the aspirin pills was ASA, which is normally not true. This assumption skewed our results by a large amount, once again making our results unreliable and

We made sure the solution is strongly acidic by testing it with litmus paper getting a pH of 2. We then cooled the mixture to room temperature swirling the flask occasionally in an ice bath. We collected the aspirin by vacuum filtration and washed the aspirin on the filter with cold distilled water. We let it air dry for 30-35 minutes and then weighed the aspirin. It weighed out at 0.513g. The unknown component was calculated and weighed at 0.738g.

Abstract: This experiment utilizes skills of titration in order to find the percent mass of actual aspirin (acetylsalicylic acid) in store-bought aspirin. Two steps and two kinds of titrations are used. First, approximately 0.1 M NaOH solution was standardized using titration with KHP. KHP was used because it does not absorb moisture when it is stored, and it is soluble in water. The reaction that takes place in this reaction is a 1:1 mole ratio reaction. The titration determines the volume of NaOH solution needed to completely react with the known amount of KHP. This makes the calculation of molarity of NaOH solution fairly easy. The average molarity calculated was 0.09897 M. The final part of the experiment is a titration to indicate the amount of acetylsalicylic acid present in approximately 0.5 g of store-bought aspirin. The volume of 0.09897 M NaOH solution indicated the moles of NaOH that reacted with the aspirin. The reaction between NaOH and acetylsalicylic acid is a 1:1 mole ratio reaction. The calculated mass of acetylsalicylic acid was 0.3963g. This was divided by the initial amount of aspirin measured (0.4976g), and multiplied by 100 to give the percent mass. This value is 79.66%. This particular experiment may not be very accurate. Specifically, only two standardization trials were counted. This may cause a less accurate measurement of NaOH solution molarity, ultimately leading to an inaccurate percent mass of acetylsalicylic acid.

This experiment involved three steps: synthesis of aspirin, isolation and purification, and the estimation of purity of the final product. The synthesis involved the reaction of salicylic acid and acetic anhydride in the presence of a catalyst, phosphoric acid, H3PO4. When the aspirin was prepared, it was isolated and filtered. The percentage yield of the synthesis was calculated to be 78.42%. The experimental melting point range of aspirin was determined to be 122 -132°C. Due to its wide range, and lower value than that of the theoretical melting point of 136°C, it was

Therefore, the experimental yield is larger than theoretical yield. This indicates an error in how the result was produced, this could have been caused by a number of different factors. The first possible reason for this error could be a result of the drying process of the amino group. The product was left out for several minutes, however despite this the product was still wet after the drying process over should have been complete, thus increasing the mass of our product and therefor increasing the percent yield. Another possible factor in this over calculation is an excess of 4-aminophenol, which was 1.008g instead of 1 gram. Additionally, a number of errors could have been the result of time management, which in the future will need to improve. In the process of cooling reactants (4-aminophenol and acetic acid) in ice water, time could have been used more effectively in other parts of the experiment. This could have been a major factor in why 3 hours were used to complete the lab. Despite these potential errors we found that both acetaminophen and 4-aminophenol were alcohols as they produced a color change in the Cerium Test. This means that both these compounds have a hydroxyl group which coincides with the resulting chemical structure record on our green sheet. Using the Universal Indicator Test the pH of 4-aminophenol was found to be slightly basic whereas the acetaminophen was acidic. These