[Chemical Engineering Laboratory I]
SEGi University ANALYSIS OF ASPIRIN Candidate’s Name: Kee Name: Kee Tze San Student ID: SUKD1601886 Group Number: 2 Group Member’s Name: 1. Chow Yee Lin
SUKD1502359
2. Tan Xin Ren
SUKD1500882
3. Tan Jee Yuen
SUKD1600614
Lecturer/ Supervisor: Ms. Nazlina bt Zulbaldi Date of Experiment: 10.10.2016 Date of Submission: 17.10.2016
1.0 Abstract
The objective of this experiment is to analyze anal yze the purity of aspirin by acid-base titration. The purity of aspirin or acetylsalicylic acid can be analyzed by using aacid-base cid-base titration. This can be done by weighing 0.5g of the aspirin prepared in the previous experiment into a clean Erlenmeyer flask. 25ml of alcohol is then added into the flask to dissolve the aspirin and two drop of phenolphthalein is added into the flask as an indicator to the color change happened during the titration. 0.1M of NaOH was then used to titrate titrate the sample to a faint pink end point. The volume of base used is then recorded. The volume of base used is the volume required to neutralize all acids in the sample of aspirin prepared. This includes impurities and acetylsalicylic acid present in the sample. An addition of 15ml of base is added to the titrated solution, this volume of NaOH is added to the Erlenmeyer flask from the burette. The mixture is heated in water bath for 15 minutes at the temperature 90°C to 95°C. Add two more drops of phenolphthalein if the solution is not pink visible anymore. For back-titration, HCl was used to titrate the excess base until the pink colors disappear. The initial and final volume of HCl u sed for back-titration is then recorded. 2.0 Theory/Introduction
In the previous experiment, we’ve learned how to prepare aspirin. The aspirin prepared is most likely not 100% pure. Even most commercial aspirin tablets are not 100% pure acetylsalicylic acid. This is because most aspirin tablet is bind by a type of chemically chemicall y inert binder to hold it in a form of tablet. One of the chemical properties of aspirin is that it is slightly slightly soluble in water because of the anhydride group present in chemical compound. Therefore, when exposed to humid air or moisture, it will hydrolyze and decomposed slowly in the surrounding. The product formed from the decomposition of aspirin is ethanoic acid. Ethanoic acid is responsible for the pungent smell present in vinegar. OH
OH H+ catalyst O O O
+
H2O
O OH
+
O OH
In this experiment, the purity of aspirin prepared will be determined by percentage. Back titration would be utilized to analyze the purity of the aspirin prepared. Normally, titration is used to determine determine the concentration of an unknown acid or base. First, titration is carried out to determine the amount of base requ ired to neutralize all acid material present in the aspirin. The solution is then added with excess base to further hydrolyze the solution. The hydrolyzed solution is then back titrate using acid. By B y using this technique, the amount of reactants not reacted in the excess can be determine. Hence, the number of milimoles of reagent reacted can be calculated by using this formula: Milimoles reagent reacted = total milimoles – milimoles – milimoles milimoles back-titrated At low temperature, aspirin can be neutralized with base:
HO
OO-
O O
O
+
+
OH-
H2O
O O
If no impurities are found in the compound, compoun d, a normal acid-base titration can be b e used to determine the purity of the aspirin. However, if imp urities are present, the base will neutralize both the acid present in the acetylsalicylic acid and the acid present in the impurities. Thus, the volume of the base required to neutralize the solution can be used to determine the number of milimoles present in the acid. Total milimoles acid = milliliters NaOH x molarity NaOH At high temperature, aspirin can be neutralized with excess base:
OO-
O
+
+
OHHO
O
OO
CH3CO2-
To determine the amount of acetylsalicylic acid present in the aspirin prepared, excess base is added to the titrated solution and placed in hot water bath to speed up the reaction. This process is known as saponification of esters. In high temperature temperature and with the present of excess base, the ester will undergo hydrolysis forming alcohol and carboxylic acid. The excess based not reacted in the hydrolysis will be determined b y a back-titration with acid HCl. With all the data collected, we can now no w determine the mass of acetylsalicylic acid present in the aspirin we prepared.
3.0 Apparatus and Material
Apparatus 1. 250ml Erlenmeyer Flask (3) 2. 50ml burettes (2) 3. 600ml beaker 4. Burette clamp 5. Water bath
Materials 1. Aspirin 2. Phenolphthalein 3. 95% ethyl alcohol 4. 0.1M HCl 5. 0.1M NaOH
4.0 Procedures
1. 5.0g of the aspirin prepared in the previous experiment was weigh into a clean, cl ean, dry 250ml Erlenmeyer flask. 2. 25ml of ethyl alcohol was added to the flask and the flask was swirled to dissolve the aspirin. Two drops of phenolphthalein were added. 3. The sample with 0.1M of NaOH was titrated to a faint pink end point. The volume of NaOH used was recorded. This volume of base corresponds to that which is required to neutralize all acids present in the sample, impurities as well as the acetylsalicylic acid.
4. 15ml of the volume of base required in the previous titration was added. Abou t this volume of NaOH was added to the Erlenmeyer flask from the burette. 5. The mixture was heated in a water bath at temperature 90°C to 95°C for 15 minutes to hydrolyze the aspirin. The flask was swirled occasionally. 6. The flask was cooled to room temperature by b y running it with cold tap water. Two more drops of phenolphthalein should be added if the solution is not pink. 7. The initial volume of HCl was recorded and the excess base was titrated using HCl until the pink color disappears. The volume of HCl used was recorded. 5.0 Results
Molarity of NaOH
=___
0.1 _____M
Molarity of HCl
=_____0.1_______M
Mass of aspirin
=_____0.5_______g
Volume of NaOH required to neutralize all acid material Final reading
=____
57______ml
Initial reading
=____ _ 20______ml
Volume of NaOH
=______37______ml
Milimoles of NaOH
=______3.7_______
Volume of NaOH used in hydrolysis Final reading
=______89______ml
Initial reading
=_____ 37______ml
Volume of NaOH
=______52 _____ml
Milimoles of NaOH
=______5.2______ml
Volume of HCl in back titration Final reading
=______57______ml
Initial reading
=______ 0______ml
Volume of HCl
=______57______ml
Milimoles of HCl
=______5.7_______
Milimoles of acetylsalicylic acid
=_____-0.5______
Grams of acetylsalicylic acid in sample
=_____-0.09_____
Purity of aspirin (%)
=_____-18_______
5.1 Sample calculation
Calculation of milimoles of acetylsalicylic acid Milimoles of acetylsalicylic acid= milimoles of NaOH used in hydrolysis - milimoles of HCl used in back titration 5.2 - 5.7 = -0.5
Calculation of grams of acetylsalicylic acid in sample Grams of acetylsalicylic acid in sample molecular weight of acetylsalicylic acid x moles of acetylsalicylic acid =
1000 180 x -0.5 =
1000 =
Calculation of purity of acetylsalicylic acid
-0.09
Purity=
grams of acetylsalicylic acid mass of aspirin
=
-0.09 0.5
x
100%
x100%
= -18%
6.0 Discussion
The aim of this experiment is to investigate inv estigate the purity of the aspirin made during the previous experiment. The techniques used to determine the purity of the aspirin made were titration and back-titration. The purpose of doing titration is to neutralize all acid present in the aspirin and the back titration is used to determine the milimoles of acetylsalicylic acid present in the aspirin. At the end of the ex periment, we’ve noticed that the percentage purity of the aspirin we made was -18%. The obtained is definitely not the ideal result. The reason and error behind the result obtained will be further discussed as below.
Figure 1: Before titration: solution is colorless
Figure 2: After titration: colorless to light pink
The mlimoles used to titrate the acid present in aspirin was 3.7 milimoles whereas the milimoles used in the hydrolysis process were 5.2 milimoles. On the other hand, the milimoles used in HCl for back-titration was 5.7 milimoles. By the end of the titration and back-titration, the solution should experience a color change of from pink to colorless.
Figure 3: Before back-titration: 2 drops of phenolphthalein were added to increase the intensity
of pink.
Figure 4: After back-titration: solution turns from pink to to colorless
From the experiment, the purity level of the aspirin falls into the negative percentage range. The aspirin prepared might not be prepared in a proper wa y as we did not recrystallize the sample of aspirin we’ve prepared in the previous experiment. Thus, causing some impurities present in the aspirin sample. The back-titration technique used in this experiment involves the titration of excess base with acid. Base used in this experiment was NaOH and the acid used was the common acid HCl. The amount of milimoles reacted from the hydrolyzed ex cess base with the acid will give us the number of milimoles of acetylsalicylic acid present in the solution. The outcome of this experiment was far from ideal be cause of the errors made during this experiment. There are high chances chanc es that the molarity prepared for the titration and back-titration is incorrect. Before any titration is carried out, we alwa ys have to make sure the molarity of the bases is correct. The failure to obtained 100% purity of acetylsalicylic acid is also due to the impurities present in the aspirin we prepared. Therefore, we n eed to carry out recrystallization when preparing aspirin. Lastly the data obtained was inaccurate because there is no duplication du plication of data. Thus, we should repeat the experiment for at least 3 times to obtain a better average of the results.
7.0 Conclusion
In conclusion, the experiment was not successful because the results obtained deviates far from the ideal results we expected. Percentage purity of the aspirin sample is -18%. The negative value portrays the molarity of base used in titration is not the correct molarity or either there is too much of impurities present in the sample. However, we have carried out the titration and back-titration successfully. The color changes in the solution indicates there is an reaction occurred with the base and acid used for titration.
8.0 References
1. Chemguide.co.uk. (2016). acid anhydrides and water, alcohols or phenol. [online] Available at: http://www.chemguide.co.uk/organicprops/anhydrides/oxygen.html [Accessed 16 Oct. 2016]. 2. Aspirin - Should it be The Ideal Medicine?. Medicine? . (2016). Chemical Properties. [online] Available at: http://aspirinpositivenegativeeffects.weebly.com/chemical-properties.html [Accessed 16 Oct. 2016]. 3. Determination of Aspirin using Back Titration. (2016). 1st ed. [ebook] Available at: http://www.lasalle.edu/~prushan/BackTitration-lab4.pdf http://www.lasalle.edu/~prushan/BackTitra tion-lab4.pdf [Accessed 13 Oct. 2016]. 20 16]. 4. Chemguide.co.uk. (2016). hydrolysis of esters. esters. [online] Available at: http://www.chemguide.co.uk/organicprops/esters/hydrolysis.html http://www.chemguide.co.uk/organicprops/esters/hydrolysi s.html [Accessed 16 Oct. 2016].
