• Join over 1.2 million students every month
  • Accelerate your learning by 29%
  • Unlimited access from just £6.99 per month

The objective of this experiment is to calculate the % purity of aspirin, by first neutralizing the acetylsalicylic acid with excess sodium hydroxide, and then back titrate the excess

Extracts from this document...

Introduction

The titration of aspirin Objective : The objective of this experiment is to calculate the % purity of aspirin, by first neutralizing the acetylsalicylic acid with excess sodium hydroxide, and then back titrate the excess base with standardised hydrochloric acid. Apparatus : burette, pipette, volumetric flask, funnel. Procedure : 1) About 1.5 gm of aspirin was accurately weighed, the mass was recorded. 2) The weighed amount of aspirin was transferred into a beaker , 25 cm3 of 1 molar NaOH was then transferred into the beaker. The beaker was then gently heated for 10 minutes. 3) After heating, the solution was then transferred into a volumetric flask via a funnel. More distilled water was then added to the flask to make a final volume of 250 a cm3. ...read more.

Middle

7) The above titration is repeated 3 times, and the result is recorded. Result and Calculation: Mass of aspirin = 1.504 gm Concentration of NaOH = 1 mole/L Concentration of HCl = 0.1 mole/L Reaction after adding of NaOH is : 2H2O 1 mole of aspirin will react with 2 moles of NaOH Reaction for the titration is : NaOH + HCl ==> NaCl + H2O 1 mole of NaOH will react with 1 mole of HCl The result of the titration is tabulated as follows: Burret Reading Titration 1 Titration2 Titration Final Reading 9.2 16 25 Initioal reading 0 7.0 16 Volume used 9.2 9 9 % purity 94.55% 95.74% 95.74% Calculation : Sample calculation using result of titration 2 # moles of HCl use = Molarity * Volume in litres ...read more.

Conclusion

Of the above only 0.09 moles was left behind, so the aspirin has reacted with 0.025 - 0.009 = 0.016 moles of NaOH. 1 mole of aspirin will react with 2 moles of NaOH So the number of moles of aspirin must be 0.016/2 = 0.008 moles Molecular mass of aspirin (C9H8O4) = 180, so mass of acetylsalicylic acid = 0.008 * 180 = 1.44 g So % purity of aspirin = (1.44/1.504) * 100 = 95.74% The other titration results was calculated the same way. Average purity = 1/3 (94.55 + 95.74 + 95.74) = 95.34% Conclusion and discussion: The error in this lab is in the judgement of the end-point, in which the solution will completely becomes colourless. The result was quite consistent, the average % purity is 95.34% ...read more.

The above preview is unformatted text

This student written piece of work is one of many that can be found in our GCSE Aqueous Chemistry section.

Found what you're looking for?

  • Start learning 29% faster today
  • 150,000+ documents available
  • Just £6.99 a month

Not the one? Search for your essay title...
  • Join over 1.2 million students every month
  • Accelerate your learning by 29%
  • Unlimited access from just £6.99 per month

See related essaysSee related essays

Related GCSE Aqueous Chemistry essays

  1. Marked by a teacher

    ANALYSIS OF ASPIRIN BY BACK TITRATION

    4 star(s)

    Hydrochloric acid (0. l mol/dm3) Aspirin tablet Bromocresol Purple indicator Deionised water A commercial aspirin tablet was placed into a 250cm3 conical flask with a magnetic stirrer bar Then sodium hydroxide solution (0.1mol/dm3, 60cm3)

  2. Determining the purity of Iron Wool.

    The funnel and any used/dirty apparatus were rinsed out using deionised waster as well. Using a pipette 25cm3 of the stock solution was extracted and place in a clean conical flask with the addition of 20cm3 of 1M sulphuric acid which had been measured in a clean measuring cylinder.

  1. Titrating Sodium hydroxide with an unknown molarity, against hydrochloric acid to find its' molarity.

    Care should also be taken to avoid inhaling any fumes that may be released by the reacting chemicals as they may cause damage to the respiratory system. Eye protection must be worn as contact of these chemicals with the eyes could cause severe damage.

  2. Making Aspirin.

    The researchers proposed that additional polymorphs of aspirin might be found if experimental crystallization conditions could be developed that would stabilize the planar conformation. Such ideas will help researchers in their attempts to find additional experimental forms of aspirin (Reference: Payne, R.C. Rowe, R.J. Roberts, M.H. Charlton, R. Docherty, J.

  1. Preparation of aspirin - The chemical background This is the overall reaction that ...

    it is reactive but not too unpleasant or dangerous, it is cheaper, less corrosive and its reaction with water is slow, so it is not readily hydrolysed. A much more reactive ethanoylating (acyl) agent is ethanoyl chloride: Ethanoic anhydride ?both acid derivatives?

  2. Experiment to produce acetylsalicylic acid (Aspirin).

    Salicylic acid Aspirin is usually prepared by reacting salicylic acid with acetic anhydride in the presence of concentrated sulfuric acid as a catalyst. This reaction is known as an Ester Formation reaction. Hypothesis to be tested (Objective) Does the Aspirin produced in the laboratory have a different melting point and

  1. Investigating different types of aspirin and making aspirin.

    With stirring until the volume is three times bigger. You must take great caution in handling the apparatus because it is extremely hot. Thin layer chromatography (TLC) Chromatography can be used to separate a mixture of components. Both qualitative and quantitative data can be obtained.

  2. the synthesis of azo dyes, aspirin and soap

    At higher temperatures above 10�C, there is too much so therefore the diazonium ion gives way to nitrogen gas. Below shows the reaction mechanism of for the formation of diazonium salt. Now that there is a N2+ diazonium ion, there is now a suitable electrophile to attack a phenol ring.

  • Over 160,000 pieces
    of student written work
  • Annotated by
    experienced teachers
  • Ideas and feedback to
    improve your own work