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  Introduction Aspirin is commonly used as an analgesic for any types of pain, an antipyretic to reduce fever or anti-inflammatory medication (American Society of Health-System Pharmacists, 2016). The drug is commonly used in several types of conditions such as in treating special inflammatory conditions, decreasing the risk of certain types of cancer, and preventing the occurrence of heart attacks (Patrignani et al., 2016). The drug is usually in the form of a tablet, which is sometimes combined with other drugs such as acetaminophen, caffeine, and other substances (Britannica, 2019). The precursor of aspirin was first extracted by the native Americans from willow and poplar barks around 2, 500 years ago. The extraction was made into a tea, which was then used to reduce fever (Ravina, 2011). In 1763, Reverend Edward isolated and identified one of the compounds that synthesize aspirin, which is salicylic acid. The compound is made available for everyone in large quantities, which later on causes severe stomach irritation to the users (Jones, 2005). Thus, in 1893, Felix Hoffman derived a derivative of salicylic acid called acetylsalicylic acid, which is less harmful to the mucous lining of the stomach. Each tablet is cloaked by an acetyl group, which is an acid-resistant envelope that passes through the stomach and small intestine. The tablet is then converted back to salicylic acid, the compound that enters the bloodstream (Jeffreys, 2008). Acetylsalicylic acid, which was produced synthetically, was given the trade name, aspirin. Aspirin is the trade name for the organic compound named acetylsalicylic acid (ASA, CH3COOC6H4COOH or C9H8O4). The drug is an acetyl derivative of salicylic acid, which is called ortho-hydroxybenzoic acid (, n.d.). It is a cyclic compound composed of a benzene ring along with 2 functional groups: the carboxylic acid (-COOH) and ester (-OCOCH3). The molecule is stable in dry air however, it gradually hydrolyzes with moisture to acetic and salicylic acid. Consumption of aspirin might result in an upset stomach, which is due to the presence of an acid functional group which can irritate the stomach lining (Stubbings, n.d.). Thus, intake of aspirin must be in moderation as it may damage the mucous membranes of the intestines, which may cause bleeding. Fig. 1.  Aspirin ASA is a monoprotic weak acid that can be prepared by reacting salicylic acid and acetic anhydride in the  presence of an acid catalyst, either sulfuric acid or phosphoric acid. It is not quite soluble, however, in the synthesis of ASA, it will precipitate once water is added while the excess acetic acid will dissolve in water. salicylic acid + acetic anhydride → aspirin + acetic acid  C 7 H 6 O 3(s)  + C 4 H 6 O 3(l)   → C 9 H 8 O 4(s)  + C 2 H 4 O 2(aq)   Fig. 2. Synthesis of acetic acid from salicylic acid and acetic anhydride  Figure 2 shows the synthesis of acetic acid from salicylic acid and acetic anhydride, which is a form of esterification reaction. The compound is a derivative of salicylic acid (Williamson et al., 2010). Salicylic acid contains two functional groups: a carboxylic acid and a phenol group, which causes stomach irritation. An ester is formed from the phenol group and carboxylic acid on the acetic acid. The reaction is desirable to replace one of the acid groups, thus, reducing its acid strength, making it easier to digest and less irritable in the stomach (Esobel, 2011). Tablets may contain other components aside from its active ingredient. These inactive ingredients may help in the production of consistent product for the consumers. Thus, in analyzing the composition of an aspirin tablet, it was hydrolyzed by an alkali, which then results in the two-component acids of the substance (austriajamesonphysics102lab, 2013). The process is called the neutralization process. An acid, ASA, reacts with a base, NaOH to create salicylate and acetate. acid + base → salt + water   C 9 H 8 O 4(s)  + NaOH (aq)   → C 9 H 7 0 4(s)  Na + H 2 0 (l)  C 9 H 8 O 4(s)  + NaOH (aq)   → C 7 H 5 0 3(s)-  Na +  + CH 3 COO -  Na + + H 2 0 (l)  The neutralization reaction can be used to determine the amount of aspirin (acetylsalicylic acid) present in commercially available aspirin tablets using a back or indirect titration method. In a direct titration setup, a standard titrant is added to the analyte until the endpoint is reached. On the other hand, in back titration is a titration method that determines the concentration of the analyte by reacting it with a certain amount of excess reagent. The remaining excess reagent is titrated with another reagent that shows the amount of excess reagent was used in the first titration (Helmenstine, 2019). The back-titration method is used whenever the molar concentration of the excess reactant is known while the concentration or strength of the analyte is unknown. Back titration is used as a method if one of the reactants is volatile, the analyte is an insoluble salt, a reaction is too slow, or the reaction involves a weak acid and weak base (Ernest, 2015). Aspirin is a weak acid that undergoes slow hydrolysis; thus, the reaction with NaOH, a strong base, is slow, in which direct titration is unfavorable to use. The activity is a demonstration of using back titration as a method of analyzing acids and  bases. The entire activity was made to demonstrate a back-titration process in identifying the percent purity of acetylsalicylic acid in aspirin tablets. The aim of the activity was to assess the potency of the drug by reacting it with NaOH in titration. The obtained composition of the tablet (percent ASA) was compared to the known ASA content of commercialized aspirin tablets. 2. Methodology Three samples of 0.100 g of aspirin were weighed, and were added with 25.00 mL of standard 0.1 M NaOH solution. The solutions were simmered gently for around 10 minutes to quicken the hydrolysis of the acids, and to ensure the completion of the neutralization of the acid. Furthermore, a homogeneous mixture of the sample, which was ideal, was created for titration. Moreover, the solutions were then quantitatively transferred and diluted to mark in a 100-mL volumetric flask. The sample was in diluted in order to eliminate interferences from other substances that may be present in the sample, to ensure that the sample was a homogeneous mixture, and to lessen the concentration of the analyte (Hach, 2018). A 50.0 mL aliquot of the sample mixture was transferred to a flask for titration. The aliquot serves as a sub-sample that was extracted from the srcinal sample. Aliquoting was done to have lesser volume of the NaOH solution reacting for neutralization, and to also have lesser volume of the HCl solution needed in back-titration (Reid, 2018). The analyte, aspirin and NaOH solution, were then titrated using HCl as the titrant until the cloudy or white color endpoint was reached. In checking the result, a drop of the NaOH was added to the sample. A pink color indicated that the correct endpoint was achieved in the back-titration  process.  3. Results and Discussion The table below shows the results obtained from the back-titration of aspirin. The % ASA and the mass of ASA per sample were shown in the table. Table 1.  Percent ASA (acetylsalicylic acid) % ASA Mass of ASA per sample (g) Trial 1 54.3306 55.4172 Trial 2 44.7182 45.0760 Trial 3 47.0997 47.9946 Mean 48.7162 mg 49.4960 Standard Deviation 5.0059 5.3316 Confidence Intervals at 95% 48.7162 ± 12.4277 49.4960 ± 13.2362 The table shows the percent purity of ASA in every sample taken during the activity and the mass of ASA in every sample. The percent of ASA shows that most values range from 44% to 54%. On the other hand, the mass of the ASA in every sample also ranged from 45 g to 55 g, which is also quite near to the percent purity calculations since the mass is related to the percent purity as it is one of the variables that must be considered in getting the percent purity. Aspirin tablets are mainly composed of acetylsalicylic acid or ASA. Getting the percent weight per weight of ASA in an aspirin tablet will determine its percent purity. Table 1 shows the percent weight per weight of acetylsalicylic acid in a 0.1000 g of aspirin tablet sample. On average, the aspirin sample has 48.7162% (w/w) of acetylsalicylic acid (ASA) with a standard deviation of approximately 5.0059. This suggests that the aspirin tablet sample, on average, is 48.7162 percent pure and has about 49.4956 mg of ASA or acetylsalicylic acid. The data disagrees with the data of Asjali et al. (2015) where approximately 71.53% to 95.97% purity of aspirin sample of the same brand was recorded. The findings further suggest that Asjali et al. recorded around 71.53 mg to 95.97 mg of a 100 mg sample of the said brand of aspirin, 22.0344 to 46.4744 larger than the data recorded of this experiment and is far beyond the 36.2885% to 61.1439% weight per weight ASA or purity at 95% confidence. The figure shows the percent ASA in each trial with respect to the mean percent purity of aspirin, as obtained in the experiment. Fig. 3 . Percent purity of acetylsalicylic acid Figure 3 shows the plot of the recorded %ASA in the three trials. The results suggest that the recorded data is  precise, which only have a standard deviation of approximately 5.0059. The deviation can have been caused  by some random errors in the duration of the experiment. These errors are unavoidable types of error; however, it can be resolved through estimating and getting the average of the recorded measurements. Thus, it can be deduced that the measurements were precise and reliable. However, the accuracy of the data recorded has a great issue. The theoretical ASA (acetylsalicylic acid) or aspirin content of a low dose aspirin tablet of the same brand used in the experiment is around 80 mg to 81 mg, or 80% to 81% aspirin content per 100 mg sample. However, this experiment only recorded about 49.4956 mg of aspirin content per 100 mg sample, which is approximately 30.5044 mg smaller than the expected aspirin content of a low dose aspirin tablet. The cause of this inaccuracy might be traced as to how the researchers did the experiment. It is possible that there were some steps or some proper laboratory techniques that were neglected that altered the data.  Moreover, the aspirin tablet sample in the activity is has a lower percent purity than the low dose aspirin tablets that usually have 75 mg to 81 mg of aspirin content. This idea implies that the aspirin tablet sample is much weaker than the very low dose aspirin tablets which further suggests its effectivity (Nordqvist, 2017). 4. Conclusion and Recommendations The % acetylsalicylic acid (ASA), which also suggests the % purity of the aspirin tablet sample is approximately 48.7162%. The result also corresponds around 48.7162 mg of ASA in a 100 mg sample of the aspirin tablet, which is about 30.5044 mg smaller than the expected weight, 80 mg to 81 mg, of ASA in the sample. The experiment recorded 5.0059 mg standard deviation, which suggests that the data recorded were  precise and reliable. However, the accuracy of the experiment was a subject of doubt since the experiment suggests that there was about 36.2885 mg to 61.1439 mg of ASA in a 100 mg sample of an aspirin tablet, which is relatively smaller than the expected weight of ASA in a sample. The inaccuracy may have been possibly caused by the negligence of some laboratory techniques such as an improper reading of the volume in the burette, negligence of air bubbles in the burette used, or carelessness in the contamination of the chemicals and reagents used in the activity. Furthermore, it is recommended to always check if proper laboratory techniques are followed. Also, more trials must be done to truly determine the % ASA or % purity of the aspirin tablet since the more trials are made, higher chances of getting and achieving the true value. It may also be possible for future experiments to analyze another type of acid using the back-titration method. References Jeffreys, D. (2008).  Aspirin: the remarkable story of a wonder drug  . Bloomsbury Publishing USA. Jones, A. (2005). Chemistry: an introduction for medical and health sciences . John Wiley & Sons, p. 5-6. Ravina, E. (2011). The evolution of drug discovery: from traditional medicines to modern drugs . John Wiley & Sons, p. 24. Williamson, K. L., & Masters, K. M. (2010).  Macroscale and microscale organic experiments . Cengage Learning. Palleros, D. R. (2000).  Experimental organic chemistry . Wiley & Sons. austriajamesonphysics102lab. (2013, May 5). Quantitative Determination of Acetylsalicylic Acid in Aspirin Tablets by Back-Titration. Retrieved from Ernest. (2015, April 17). How does a back titration differ from a regular titration?: Socratic. Retrieved from Esobel. (2011, May 23). Full Report: Synthesis of Aspirin. Retrieved from 
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