Titration Of Aspirin Information

About Titration Of Aspirin

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A laboratory technician wants to determine the aspirin content of a headache pill by acid-base titration. Aspi A laboratory technician wants to determine the aspirin content of a headache pill by acid-base titration. Aspirin has a Ka of 3.0 x 10-4 M. If the pill is dissolved in water to give a solution about 0.005 M, what is the pH of this solution? (Neglect dilution effects.) If the solution in the problem above is then titrated against KOH solution, what will be the pH at the stoichiometric point.

cattbarf replied: "With this Ka and concentration, you can't use the usual formula for weak-acid equilibrium X^2/[Acid conc]=Ka because you can't neglect the amount of acid actually reacted to H+ and the Cation. So you have the more general equation, X^2/[0.005-X]= 3x10^-4 You can solve this by trial/error or by the general quad. eqtn method. Roughly, X=0.0027 At the stoichiometric point, all acid is dissociated, and the problem is equivalent to dissolving 0.005 M potassium salicylate in water. In that case, Salicylate- + H2O = Acid + OH- In this situation, the equilibrium expression is [Acid][OH-]/[Salicylate] = Kb Since KbKa=10x10-15, Kb= 3.3x10-11 If X is the amount of [OH-] formed at equil., then X^2/[0.005]= 3.3x10-11. X is roughly 4x10-7. From this you can compute pOH= 7-log 4= 6.4. Then since pH+pOH=14, pH=7.6."

What are the principles behind the use of back titration in the analysis of aspirin tablets? experiment on quantitative determination of acetylsalicylic acid in aspirin tablets by back titration

Paul H replied: "Aspirin is an ester which is very easily hydrolyzed. So easily that during a normal titration with NaOH, the alkaline conditions break it down leading to errors in analysis. In addition its water solubility is low. To overcome these problems, Aspirin is completely hydrolyzed to salicylic acid and acetic acid with hot, excess NaOH. The NaOH in excess is then titrated with standardized acid (HCL or H2SO4) to calculate the amount used in the reaction with Aspirin and thereby the amount of drug in the flask."

when testing the purity of aspirin, why is titration helpful? what does the titration show? is a smaller titre value or a larger titre value indicative of a purer sample? etc thanks fed i am dissolving the aspirin in ethanol and then titrating against NaOH (0.1moldm-3)

rje102 replied: "What are you titrating it against? Another method would be to use melting point apparatus. You can measure the melting point accurately and compare it to the known. It your sample melts over a range of temperatures it indicates it is not very pure!"

Alec replied: "Aspirin is an organic acid, and therefore it can be titrated against a strong base like NaOH. You know the molecular weight of acetylsalicylic acid (aspirin), and you know how many moles of acid you should have. Once you titrate it, it will show if your sample contained 100% aspirin. You will never have more than 100% aspirin, so if the titration shows that you have more acid that you have calculated - then either you did the titration inoccrectly, or you have another acidic impurity present. Most likely you will have less than the calculated/throretical amount based on the titration, and the will indicate how much of the mass you're titrating is other stuff - non-acidic impurities."

Chemical equation for aspirin titration? can someone help me write equation for aspirin being titrated by adding NaOH?

Merlin's Feline replied: "sure acetylsalicylic acid is C6H4( OOCCH3) COOH It is an o-hydroxy benzoic acid wherein the OH group is acetylated. This means it is a monoprotic acid. So in titrations with NaOH Sal-COOH + NaOH ------- Sal-COO^-1 Na+ + H2O you form thje sodium salt of salicylic acid . Since the sodium salt is a salt of the weal acid salicylic acid, the pH at equivalence point will NOT be 7 but rather in the 8ish range....so phenpphthalein should be a resonable indicator as it is distinctly pink at pH ~ 8.2-8.4 and red at 9. The equivalent weight of salicylic acid is the same as it formula weight ( 180 )"

what are the advantages of a back titration over a standard titration with NaOH for analysing aspirin? Is it because in a titration with NaOH the acid impurities react aswel? im really stuck.. thanks

ChemistryMom replied: "It's because aspirin is such a weak acid that it reacts slowly with the NaOH, making it difficult to accurately get a good endpoint in a reasonable time. With back titration, you react with an excess of NaOH (known amount), heat it to make the reaction go to completion quickly, then use HCl to determine the amount of NaOH that is remaining. This reaction will go quickly, and is much easier to measure."

I am doing a lab report involing the titration of a known mass of aspirin with standardized NaOH? Please tell me the equation for this reaction, is it one to one, because I need to use the equation to find aspirins experimental molecular mass THANKS!!

svrwxsooner replied: "HC7H6O3(aq) + OH-(aq) --> H2O(l) + C7H6O3-(aq) The hydroxide group pulls the proton off of the salicylic acid and creates water. This continues until the equivalence point is reached. After the eq. point then there is no more H+ for the OH- to react with so the pH becomes basic."

Titration of Aspirin. Whats the molar mass? Im stumped!? Did and experiment in lab on the titration of 325mg aspirin with 0.100 M of NaOH. I did this experiment twice so their fore Im looking for the average molar mass. Trail one I used 18.8 of NaOH and Trail 2 I used 18.4. I think the equation I would use is molar mass= grams/mole. I just dont know how to apply. HELP!

bernie_bph replied: "From the data you provided, calculate the number equivalents of aspirin. equivalents of aspirin = (0.1)(0.0188) = 0.00188 If you did not dilute your aspirin, below is your result: trial 1 = (325/1000)/(0.00188) = 172.872 grams/mole trial 2 = (325/1000) / 0.00184 = 176.630 grams/mole average = 174.751 grams/mole from source: aspirin has a molar mass of = 180.157 g/mol"

Is salicylic acid the main component of aspirin? (titration)? salicylic acid -->C6 H4 (OCOCH3 )COOH Is that the main component of aspirin ? I did a lab on aspirin titration, and I'm just confused how salicylic acid should be used to determine the purity of the aspirin; Thank you Thank you for your answers! Akeem: " (Remember the base ONLY reacts with the aspirin in the tablet BUT there's other substances in the tablet i.e. the impurities.) " I've been trying to understand that; how does that affect my result.. or does it affect it at all? I did use NaOH for the titration; I have calculated and found the purity; so does it means that the greater the amount of salicylic acid, higher the percentage of purity? And the final "% of purity" I have calculated.. that would be considered slightly inaccurate (slightly higher) because of the impurities.. is that what you're saying?

James replied: "Aspirin is Acetylsalicylic Acid, or ASA. So, to answer your question, Salicylic acid is the main component of aspirin, with a single acetyl group added on. It would be used in a electrophoresis or a simple chromatography test along-side a sample of aspirin to determine how pure the aspirin sample was. This works because Salicylic acid is a smaller molecule, and will travel further through the gel/filter paper than the aspirin would, since it has the additional acetyl group weighing it down. If the aspirin is not very pure, it will contain lots of salicylic acid, which would separate, and you would see two samples, one right beside the salicylic acid sample and one not quite as far down."

Akeem replied: "Yes aspirin is the main component.In the lab you most likely reacted the crush tablets with a base of known concentration(e.g. NaOH). The volume of base required to completely react with the aspirin in the tablet can be used to find out the number of moles of the aspirin in the tablet. Convert this number of moles to mass. (Remember the base ONLY reacts with the aspirinn in the tablet BUT there's other substances in the tablet i.e. the impurities.) The mass of the tablets used should have been weighed before hand also. To determine the % impurity: Mass of Aspirin/Mass of tablet used x 100"

Titration of aspirin with 0.5M NaOH? Hi, can someone help me with this chemistry lab question? Calculate the volume of 0.5 M NaOH that will be required to titrate a 0.15g sample of aspirin. So far I've taken this equation: NaOH + C9H8O4 ----> H2O + [C9H7O4]- [Na]+ and calculated moles of aspirin - 0.15g / 180.15 g = 0.0008 moles aspirin. Since the moles of NaOH and aspirin in the equation are equal, I assumed that at the end of the titration, moles NaOH = moles aspirin. This led me to do 0.0008 moles NaOH / 0.5 M = 0.16 mL - however, this answer seems too small. Can anyone point out what I did wrong? Thanks in advance! Hi, can someone help me with this chemistry lab question? Calculate the volume of 0.5 M NaOH that will be required to titrate a 0.15g sample of aspirin. So far I've taken this equation: NaOH + C9H8O4 ----> H2O + [C9H7O4]- [Na]+ and calculated moles of aspirin - 0.15g / 180.15 g = 0.0008 moles aspirin. Since the moles of NaOH and aspirin in the equation are equal, I assumed that at the end of the titration, moles NaOH = moles aspirin. This led me to do 0.0008 moles NaOH / 0.5 M = 0.16 mL - however, this answer seems too small. Can anyone point out what I did wrong? There are also 2 follow-up questions: 2. A student obtained the end point of the titration 0.76 mL prior to the calculated value in 1 (answer to question 1 =1.7mL). What was the purity of the sample? 3. A student required 0.40 mL more NaOH than she calculated she would need in question one. Explain how this might happen. Thanks in advance!

Dr.A replied: "0.00083 mol / 0.5 M = 0.0017 L => 1.7 mL"

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