(I) Object of the experiment
To determine the percentage of acetylsalicyclic acid in one commercial aspirin tablet, using the principle of back titration.
(II) Discussion
Back titration rather than direct titration was used because there was no suitable indicator for acetylsalicyclic acid and sodium hydroxide solution. Therefore, excess amount of NaOH was used to react with acetylsalicyclic acid. The number of moles of unreacted NaOH was determined from titration with hydrochloric acid. Eventually the number of moles of acetylsalicyclic acid and its percentage in the commercial aspirin tablet were deduced.
(III) Procedure
1. The weight of one tablet of commercial aspirin tablet was measured.
2. The aspirin tablet was grinded into powder by mortar and pestle. Then it was rinsed thoroughly with deionized water and poured into a beaker.
(Experimental set up of grinding an aspirin tablet)
3. 25cm3 of 0.2521M NaOH solution, which was in excess so that the equilibrium was driven towards products according to Le Chatelier’s principle, was pipetted to the beaker to hydrolyze acetylsalicyclic acid.
4. The solution was heated using a Bunsen burner for 5 minutes to ensure all acetylsalicyclic acid has undergone hydrolysis and fasten the rate.
(Experimental set up of heating acetylsalicyclic acid with sodium hydroxide)
5. The solution was poured into a 250.0cm3 volumetric flask. Deionized water was used to rinse the conical flask so that all the solution was transferred.
6. 0.1003 M of HCl was diluted tenfold to 0.01003 M. 25cm3 of HCl was pipetted to another 250.0cm3 volumetric flask, then deionized water was added to the volumetric flask until the graduation mark.
7. After the dilution in...
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...ch as random parallax error and the ambiguity in distinguishing between pale pink and colourless.
Random errors were reduced by taking average of repeated readings. In the titration, the amount of HCl used was almost the same among four runs. Therefore, the random error was minimized.
4. Side reactions might occur between unknown chemicals in the aspirin tablet and NaOH solution. Thus, the amount of NaOH reacted with acetylsalicyclic acid would be overcounted.
Further exploration of contents of aspirin would be needed to see whether there was any side reaction.
(IX) Conclusion
The commercial aspirin tablet had 93.24% purity. This showed that there were other chemicals like binder was presence in the tablet. Besides, the tablet might not be stored in dry condition, so acetylsalicyclic acid would be hydrolyzed by moisture in the air to form acetic acid.
The purpose of this experiment was to learn and preform an acid-base extraction technique to separate organic compounds successfully and obtaining amounts of each component in the mixture. In this experiment, the separation will be done by separatory funnel preforming on two liquids that are immiscible from two layers when added together. The individual components of Phensuprin (Acetylsalicylic acid, Acetanilide, and Sucrose as a filler) was separated based upon their solubility and reactivity, and the amount of each component in the mixture was obtained. Also, the purity of each component will be determined by the melting point of the component.
We then took 1ml of the 0.1% solution from test tube 2 using the glucose pipette and added it to test tube 3, we then used the H2O pipette and added 9ml of H2O into test tube 3 creating 10ml of 0.01% solution.
Once the mixture had been completely dissolved, the solution was transferred to a separatory funnel. The solution was then extracted twice using 5.0 mL of 1 M
acid*1 mol s. Acid/ 138.1g s acid*1 mol aspirin/1 mol s. acid * 180.2 g aspirin/1 mol aspirin = 3.9145 aspirin
We were then to make a base solution of 0.7 M NaOH. In order to standardize
Aspirin contains the substance acetylsalicylic acid (ASA), which can relieve inflammation, fever, pain, and known as a “blood thinner”. Aspirin was not officially trademarked until March 6, 1899 when the Imperial Office of Berlin made it official. It has been used for the last 110 years, but its natural form, salicylic acid has been around for thousands by Egyptians, Greeks, and Romans. Aspirin is available in over 80 countries and known as the best non-prescription drug. The most common use of aspirin is to cure headaches and use it as a pain reliever, but aspirin is known to prevent heart attack and strokes. It was first proposed in 1940, but wasn’t confirmed until 1970 when doctors would recommend taking aspirin daily [1].
7. Using the stirring wire, stir the mixture until the solute completely dissolves. Turn the heat source off, and allow the solution to cool.
Some people alternate the use of other OTC such as Aspirin, which also has other dangerous effects , but hopefully will reduce acetaminophen toxicity.
The conical vial was placed in a small beaker and allowed to cool to room temperature. The mixture was Cooled thoroughly in an ice bath for 15-20 minutes and crystals collected by vacuum filtration on a Hirsch funnel. The vial was rinsed with about 5 mL of ice water and transferred into to the Hirsch funnel and again washed with two additional 5mL portions of ice water. Crystals were dried for 5-10 minutes by allowing air to be drawn through them while they remained on the Hirsch funnel. The product was transferred to a watch glass plate and allow the crystals to dry in air. Crude acetaminophen product was weighed and set aside a small sample for a melting point determination and a color comparison after the next step. Calculation of the percentage yield of crude acetaminophen (MW = 151.2). was done and recorded in the lab notebook.
2nd step heat the mixture: Make sure the agarose dissolves. Wait until it boils and when you are going to transfer the mixture, wear gloves to avoid getting burnt. Transfer the mixture into a removable gel tray.
Firstly, we need to keep the chemical at a constant concentration. So, in this experiment we have chosen to keep hydrochloric acid at a constant concentration (5cm3). We could have, however, used Sodium Thiosulphate as a constant, but we had chosen to use Hydrochloric acid. Next, we must make sure that the solution is kept at a constant volume throughout the experiment. If the volume is different, then it could give different results if it was at a constant volume.
Analysis of Aspirin Tablets Aim --- To discover the percentage of acetylsalicylic acid in a sample of aspirin tablets. ----------------------------------------------------------------- In order to do this, the amount of moles that react with the sodium hydroxide must be known. This is achieved by using the method of back titration.
second test tube also add 6 mL of 0.1M HCl. Make a solution of 0.165
Firstly, an amount of 40.90 g of NaCl was weighed using electronic balance (Adventurer™, Ohaus) and later was placed in a 500 ml beaker. Then, 6.05 g of Tris base, followed by 10.00 g of CTAB and 3.70 g of EDTA were added into the beaker. After that, 400 ml of sterilized distilled water, sdH2O was poured into the beaker to dissolve the substances. Then, the solution was stirred using the magnetic stirrer until the solution become crystal clear for about 3 hours on a hotplate stirrer (Lab Tech® LMS-1003). After the solution become clear, it was cool down to room temperature. Later, the solution was poured into 500 ml sterilized bottle. The bottle then was fully wrapped with aluminium foil to avoid from light. Next, 1 mL of 2-mercaptoethanol-β-mercapto was added into fully covered bottle. Lastly, the volume of the solution in the bottle was added with sdH2O until it reaches 500 ml. The bottle was labelled accordingly and was stored on chemical working bench.
In this experiment three different equations were used and they are the Stoichiometry of Titration Reaction, Converting mL to L, and Calculating the Molarity of NaOH and HCl (Lab Guide pg. 142 and 143).