The Mass of Acetylsalicylic Acid in Aspirin ————————————————- Purpose The purpose of the lab is to determine the mass of the ‘active ingredient’ in a commercial ASA tablet. ————————————————- ————————————————- Introduction There are three main theories surrounding acids and bases including the Arrhenius, Bronsted-Lowry, and Lewis theories. The Arrhenius theory of acids and bases states that acids produce hydrogen ions (H+) in solution while bases produce hydroxide ions (OH-) in solution.
Arrhenius was able to determine his theory based on his prior knowledge of the behaviour of substances in aqueous solution. Arrhenius went further to describe strong acids as a strong electrolyte that was able to ionize completely in order to give hydrogen ions in aqueous solutions. However weak acids can only ionize partially, remaining moderately in molecular form. Similarly, a strong base is also a strong electrolyte that ionizes completely to release hydroxide ions in aqueous solution while weak bases only partially ionize.
The Arrhenius theory is specifically deals with acid-base reactions in water. The second theory, intends the contemporary ‘protonic’ or Brownsted-Lowry theory of acid-base behaviour. This theory states that an acid is a compound or ion that can give up a proton however a base is a compound or ion that accepts a proton. Water is amphiprotic because water can give up and accept proton depending on the reaction. Water gives up a proton to form hydroxide ions (OH-) but water accepts a proton to form hydronium ion (H3O+).
The Bronsted-Lowry theory is just an addition to the Arrhenius theory in term of all Arrhenius bases that are sources of hydroxide will accept protons. Bronsted-Lowry theory also coincides with the ability for ammonia and amines to accept protons to form ammonium ions and it can also be applied to reactions that involve solutions that do not have water as its solvents. The last theory, the Lewis theory of acids and bases, claims that acids accept electron-pairs while a base donates electron-pairs. The Lewis theory is useful to explain organic reactions. An example of this theory is the base
H3N which has a lone pair. So according to the theory, this base will donate the electron pair to a proton when forming the acid NH4+. Titration is a method used to identify the exact concentration of an unknown solution by way of a chemical reaction. The molarity of acidic or basic solutions can be used to switch to and from moles of solutes and volumes of their solutions. Titration is performed when solution 1 is added to solution 2 until a chemical reaction between the solutions is visible. Solution 1 is called the titrant and solution 1 is said to titrate solution 2 at the completion of the reaction.
The endpoint of titration is determined by a change of color which is a result of indicator added to solution 2. Assuming solution 2 is the acid, as long as there are excess H+ ions in the solution, the solution remain acidic. The indicator also stays mostly in acid form so the acid-indicator solution is colorless. However when enough base solution is added to react with the H+ ions the reaction is complete, the endpoint has been reached. However if even a drop more of base is added into the acid, there is now excess hydroxide ions in the solution.
The OH- ions of the base that has been added to the acid reacts with the indicator changing the indicator from acid to base. When using phenolphthalein as indicator, for example, the base form for this indicator is red. In this case when the indicator changed into base form, the solution turns red indicating the reaction is complete. Observing the amount of solution 1 used to titrate solution 2 and the knowledge of the concentration of solution 1 and the volume of solution 2 used, the concentration of solution can be identified. Neutralization reaction is the acid and base reaction resulting in salt and water.
The state of acid base reactions when same amount of acid reacts with same amount of base to produce equal amounts of salt and water is known as equivalence point. The completion of the reaction marks the period in which the base and acid neutralize each other. An example of neutralization reaction is: NaOH+HCl>NaCl+H2O The base, NaOH, reacts with the acid, HCl, to produce salt (NaCl) and water (H2O). Principally, neutralization reactions are double displacement reaction. Neutral point is the condition of the acid-base reactions when the roduct of the neutralization reaction is neutral (it has a pH = 7). The salt produced from a strong acid and a strong base is pH neutral and water is already known to have a pH of 7. Since the products of a strong acid and a strong base reaction have neutral products, equivalence and neutral points are the same. However when a combination of strong and weak acid and base react the product salt will not have a pH of 7. However acid-base reactions are still referred to as neutralization. In order to make an acidic less acidic, a base can be added to solution in order to reduce the acidity of the acid.
The same can be done by adding more acid to base to make the base more acidic and less basic. The acid-base indicator used in this lab is phenolphthalein. Phenolphthalein is an organic compound which may be written as C20H14O4. Phenolphthalein is colorless in acidic solution but it is pinkish in basic solution. The transition from colourless to pinkish occurs around pH 9. Phenolphthalein does not dissolve well in water, so in order to do titrations it has to be prepared in alcohol solutions. At times, when the indicator is added to the acid, it might turn cloudy white.
This is because of a precipitate of solid phenolphthalein when high concentration is higher than the solubility product. However the cloudy white will diminish if the solution is stirred. Phenolphthalein has often been used but there have been new concerns with this indicator. It is said the phenolphthalein’s carcinogenicity resulted in its replacement with other materials. There have not been any health concerns with this indicator given that low amounts of the indicator are used. ————————————————- ————————————————- Materials * * Retort Stand Burette * Burette Clamp * Funnel * Methanol * Phenolphthalein * Distilled Water Bottle * Sheet of White Paper * 2 Beaker (250 mL) * 2 Erlenmeyer Flask (250 mL) * Sodium Hydroxide Solution * 1 Aspirin Solution (ASA) * Mortar & Pestle ————————————————- Apparatus ————————————————- ————————————————- ————————————————- ————————————————- ————————————————- ————————————————- ———————————————— ————————————————- ————————————————- ————————————————- ————————————————- ————————————————- Procedure 1. Gathered all the equipment needed to perform the lab. 2. Labelled a dry beaker for each liquid: methanol and waster (NaOH was already provided in a beaker). 3. Obtained 20 mL of methanol in the respective beaker. 4. Crushed one Aspirin tablet, using a mortar and pestle. 5.
Added the 20 mL of methanol into the mortar with the crushed Aspirin and mixed well with the pestle. 6. Poured half (10 mL) the methanol-Aspirin solution into each of the two Erlenmeyer flask. 7. Poured the sodium hydroxide solution into the burette through the funnel to the 0mL level on the burette. 8. Placed a sheet of blank paper under one of the Erlenmeyer flask and put the flask on the stage of retort stand, directly under the burette. 9. Titrated the sodium hydroxide into the Erlenmeyer flask while swirling the contents. 10. Recorded the final reading of sodium hydroxide remaining in the burette. 1. Repeated steps 8 – 10 for the second Erlenmeyer flask, noting that the initial speed is the final speed from the previous titration. 12. Disposed the chemicals as directed 13. Rinsed the burette and beakers with distilled water and placed the equipment back in its rightful place. ————————————————- Experimental Results Table 1: Volume of Sodium Hydroxide Used Reading (mL)| Trial 1| Trial 2| Final Reading| 8. 3 mL| 16. 1 mL| Initial Reading| 0 mL| 8. 3 mL| Volume Used(Final – Initial)| 8. 3 mL| 7. 8 mL| ————————————————- Analysis . The readings do not agree within±0. 2 mL. However the average of the two readings recorded are as follows Average=Reading 1+Reading 22 Average=8. 3 mL+7. 8mL2 Average=8. 05 Reading 1 = 8. 3 mL Reading 2 = 7. 8 mL ? The average of the two readings is 8. 1 mL. 2. CH3COOC6H4COOH(aq) + NaOH(aq) > CH3COOC6H4COONa(aq) + H2O(l) CH3COOC6H4COOH : NaOH 1 : 1 n=0. 00081moles MM=180. 17 g/mol m=n? MM =0. 00081 mol? 180. 17gmol =0. 1459g m=146 mg V=8. 1 mL (0. 0081 L) C=0. 1 M n=CV =0. 0081 L? 0. 1 M =0. 00081moles Molar Mass CH3COOC6H4COOH C=12. 01? 9=108. 09 H=1. 01? 8=8. 08 O=16? 4=64
MM=180. 17 ? The mass of the ASA in a tablet is 150 mg. 3. The percentage difference is calculated as follows: % difference=|experimental value-predicted value|predicted value? 100% % difference=|146mg-162. 5mg|162. 5 mg? 100% % difference=9. 98 ?There is a 10% difference in the experimental value and the predicted value. 4. The ASA tablets may be coated or buffered because for some individuals may find that eating these tablets makes their stomach upset. So the buffered tablets neutralize stomach acid by increasing the pH. There two types of coated ASA tablets, micro-coating and enteric coating.
Micro-coating refers to the film coating on the Aspirin tablet. The thin coating allows individuals to swallow the tablet with ease especially people who find it hard to swallow uncoated tablets. Enteric coating is a “delayed-release safety” coating. This coating will let the tablet easily pass through the stomach to the small intestine, right before it dissolves. Enteric coating is not beneficial for quick pain relief; it is usually used patients who take the table daily. 5. The additional ingredients in Aspirin and other medication tablets are carnauba wax, corn starch, hypromellose, powdered cellulose and triacetin.
Ingredients are usually added to offset the acid that is present in the tablet. The acid may be too strong for some stomachs to handle so extra ingredients need to be added to neutralize the tablet but not to the extent that the acid is no longer effective. ————————————————- ————————————————- Discussion In the titration, a base solution, NaOH, with a known concentration is added to the Acetylsalicylic acid, in very small amounts. In most cases titration is done to identify the concentration of the acid or base, depending on which one is the unknown solution.
In terms of this lab, the unknown solution is Acetylsalicylic acid but instead of finding the concentration of this acid, it is necessary to find its mass. This can be done by using the mole to mole ratio between the sodium hydroxide and the Acetylsalicylic acid (we can’t directly find the mole of the acid because the concentration of the acid is unknown so the base is required). However in order to find the moles of sodium hydroxide, the volume of NaOH required to reach the endpoint of titration must be identified. The endpoint of titration is the amount of NaOH needed to completely react with Acetylsalicylic acid.
As the base solution is added to the Acetylsalicylic acid, the acid becomes less and less acidic. The base has to be added until the endpoint of titration is reached because prior to the endpoint there is more acid than base in the flask so a complete reaction between base and acid won’t take place to produce water and salt. The endpoint will be reached even without the addition of the indicator, phenolphthalein but without the indicator it would be impossible for humans to observe the period in which endpoint is reached. When the indicator changes colour sodium hydroxide has titrated Acetylsalicylic acid.
As soon as the colour changes a slight permanent pink, it is important to record the burette reading. If the burette reading is taken after the solution has turned bright pink, it is past endpoint of titration. The solution is now more basic than acidic. The permanent faint pink indicates that the solution at the time is neither basic nor acidic; it has a pH of 7. After recording the endpoint burette reading it is essential to subtract the starting volume of NaOH that was in the burette because the difference is the volume of NaOH it takes to react completely with Acetylsalicylic acid.
Now, the volume and concentration of the base or sodium hydroxide is available. This will provide the amount of moles of NaOH that was required titrate the Acetylsalicylic acid. Using the mole to mole ratio, the number of moles of Acetylsalicylic acid that was used can be calculated. The moles multiplied by the molar mass will finally provide the mass of acid used in the lab. The endpoint of titration identifies the volume of NaOH needed to reach with Acetylsalicylic acid to produce water and CH3COOC6H4COONa (the salt).
If the endpoint is not reached or it is exceeded then the equation is no longer agreeing with the results, in which case the equations cannot be used to find concentration or mass of Acetylsalicylic acid. There were many experimental that were encountered during the process of completing the lab. As the tablet was grinded, there was some powder that was left behind on the mortar. There were also some left over powder in the beaker. At one point the pipette was used to transfer the methanol-acid solution from the beaker into the flask.
Some of the solution was left behind in the pipette. This may have been where some of the ‘active ingredient’ was lost. The amount of substances used in the lab may have been measured inaccurately. There was an instance when 20 mL of methanol was measured but when it came to split the methanol-acid solution there wasn’t enough to divide the solution in half into each flask. The final reading of the burette may have been inaccurate. It is hard to read the burette especially when the base lies between two lines on the burette. This changes the experimental mass of the ASA in the tablet.
I also think that the exact reading at which the endpoint was reached was not recorded because during both trials the solutions did not showcase faint pink. The second trial is definitely more reliable because it was done more carefully but it still seems to be over the endpoint. The results of the experiment are relatively close to that of the amount stated on the aspirin bottle which is 325 mg. However only have of the mass of one tablet it used because the one tablet was shared between the two trials. In that case, there was only a difference of 16. mg difference between the experimental and the accepted mass of the active ingredient in aspiring. The decrease in mass for the experimental results is acceptable because of the experimental errors that were made. All the errors showed signs of a loss in the table powder or its solution. The above calculation also showed only a 10% difference in the actual and experimental records. The accepted value seems to be more acceptable because it was recorded by the company and probably certified. It also seems more reliable that the experimental result because it is not attached with a list of experimental errors. ———————————————— Conclusion The relative molecular mass of the ‘active ingredient’ in a commercial ASA tablet was determined to be 146 mg. The predictions that I produced were not accurate. My prediction was based on the fact that companies usually exaggerate the facts on commercial products so that buyers are in favour of that particular product. However based on the lab, the experimental results were relatively close to that of the mass posed on the aspirin bottle. This demonstrated that the manufacturer’s claim is accurate. The observations from the lab show that approximately 8. mL of NaOH was needed to titrate Acetylsalicylic acid. Based on calculations involving this observation, it was determined that 146 mg of ‘active ingredient’ is present in Aspirin tablet. This experiment provided a better understanding of the purpose and use of titrations. Titration is a great method to determine concentration, mass and moles of certain acids and bases in a reaction. ————————————————- Bibliography 1. “Aspirin. ” Apuntes, Trabajos, ExA? menes, PrA? cticas y Otros documentos. N. p. , n. d. Web. 28 May 2010. <http://html. rincondelvago. com/aspirin_1. html>. 2. Neutralization reaction. ” Connexions – Sharing Knowledge and Building Communities. N. p. , n. d. Web. 28 May 2010. <http://cnx. org/content/m17138/latest/>. 3. “Phenolphthalein. ” Digipac Microcomputer Software. N. p. , n. d. Web. 28 May 2010. <http://digipac. ca/chemical/equilibrium/phenolpthalein. htm>. 4. “Titration. ” An Introduction to Chemistry – Bishop. N. p. , n. d. Web. 28 May 2010. <http://chiralpublishing. com/Bishop_Titration. htm>. 5. stander. “Acids and Bases. ” Chemistry and New Zealand. N. p. , n. d. Web. 28 May 2010. <http://www. chemistry. co. nz/acids_and_bases. htm>.