Quantitative Determination of Acetylsalicylic Acid in Aspirin Tablets by Back-titration E.M.M. Medrano1, J.M. Pasco1, M.E. Lubrin1, M. Manrique2 1Department
of Mining, Metallurgical and Materials Engineering, College of Engineering Institute of Physics, College of Science University of the Philippines, Diliman, Quezon City, Philippines Date Due: 21 January 2014 Date Submitted: 21 January 2014 2 National
RESULTS AND DISCUSSION The purity of aspirin was expected to be determined by the end of the experiment using the concept of back-titration. Back-titration involved the addition of large amounts of strong base which neutralized the acids formed . Then, the excess was back-titrated using a strong acid. [1] The amount of acid used to back-titrate is equal to the amount of excess base. Therefore to determine the amount of base that reacted with ASA (acetylsalicylic acid), the following equation is used:
mmolreactedNaOH = mmoltotalNaOH − mmolexcessNaOH (1) where mmolexcessNaOH is equal to the millimoles of HCl that was used to back-titrate. Back-titration is an indirect procedure. [2] This procedure is commonly used for slow reactions. In this experiment, the degradation of aspirin's acids was hastened by the addition of large amounts of NaOH with higher molarity than the base solution used in standardization. This, together with simmering the solution ensured that the component acids would completely dissociate and be neutralized. The solution was diluted after simmering so that the molarity of the base will approximately be equal to the standardized base solution. The aliquot was then titrated instead of the whole solution so that only the excess base would react with the acid. Solutions of sodium hydroxide are not suitable primary standards since the NaOH pellets cannot
be accurately weighed.[3] Lab results show that a supposed 0.05M NaOH solution is actually 0.0489M. A huge difference in the molarity of standardized solutions and solutions used in the analysis can be a source of error. A higher molarity of standard solution will result to higher %ASA of aspirin while a lower molarity of standard solution will result to lower %ASA. CONCLUSION The purity of aspirin was determined in the experiment using the concept of back-titration. It was found out that the aspirin sample has relatively low acetylsalicylic acid. Table 5.1 Percent ASA of Aspirin Trial 1 2 3 %ASA
28.56%
27.50%
27.99%
Average 28.02% ± 1.3%
It is important to determine the percent ASA of aspirin since it presents the aspirin's potency as a drug. An effective aspirin has about 70% ASA. The particular aspirin sample used was prepared in advance. Future studies may opt to prepare the aspirin samples right before they are analyzed to make sure that the drugs still have high effectivity. Consequently, analysis of ASA in aspirin may be done through direct titration. Instead of dissolving the aspirin sample in NaOH and simmering it, ethyl alcohol may be used to dissociate the component acids. [4] Ethyl alcohol and acetylsalicylic acids are both polar so ethyl alchol can dissolve ASA.
REFERENCES [1-2] Institute of Chemistry Analytical Chemistry Laboratory Manual, 2007 Revised edition; University of the Philippines, Diliman, Quezon City, 2007; p. 45 [3] Spurlock, D. Standardization of a Base, NaOH Class Notes. Indiana University Southeast. [Online] 2013. http://homepages.ius.edu/ DSPURLOC/c121/week11.html (accessed Jan. 20, 2014) [4] Pall, S. Analysis of Aspirin. [Online] 2010. http://www.savitapall.com/medicines_drugs/l abs/Analysis%20of%20Aspirin.pdf (accessed Jan. 20, 2014)
