The Purity Of An Aspirin Sample

The aim of my investigation is to find the most accurate method of measuring the purity of the aspirin through the use of quantitative analysis methods, including(1):

– Acid-base titration

– Back titration

– Colorimetry

The most likely impurities present in an aspirin sample are salicylic acid (un-reacted or as a result of hydrolysis) or ethanoic acid (or other acid catalysts).

Introductrion to aspirin

What is aspirin?

Aspirin (2-ethanoylbenzoic acid or Acetylsalicylic acid) is probably the most commonly used pain relief medicine in the world today and is described as having analgesic (pain-killing), anti-inflammatory and antipyretic (fever-reducing) actions. The calcium salt is now marketed as a soluble solution, and the sodium salt as effervescent (fizzy) aspirin. 4000 million aspirin tablets are sold in the UK every year(2). Pure aspirin is a white, crystalline powder that is synthesised artificially in the reaction:

(3)

The discovery and artificial synthesis of aspirin

The use of willow bark and leaves for pain relief dates abck to 400BC, when Hippocrates recommended it for easing the paing of childbirth. In 1763, Reverend Edmund Stone used a will bark brew to reduce fevers. The useful chemical in willow bark is salicin, which has no pharmocological effect by itself. However, when salicin is ingested it is converted through oxidation and hydrolysis to salicylic acid (or 2-hydroxybenzoic acid), and active chemical.

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In order to synthetically synthesise salicylic acid in a Labratory a chemical of similar structure can be converted through a simple chemical reaction. The structure of phenol only differs from salicylic acid by one functional group, a carboxylic acid group. Initially salicyclic acid was marketed by Bayer as a fever cutting pain relief medcine, however the bad side effects of using salicylic acid became known not long after.

Felix Hoffman was the first person to synthesise aspirin (2-ethanoylhydroxybenzoic acid or acetylsalicylic acid) in 1897. He did this by altering the structure of salicylic acid and testing the effectiveness of the new compound on his father, who suffered from chronic rheumatism. Hoffman successfully prepared a derivitive of salicylic acid that was as effective as effective as pure salycilic acid but without the unpleasant side effects. The use of acetylsalicylic acid (aspirin) as pain relief began. Aspirin was initially sold in sachets but Bayer deicded to pellet the powder and aspirin became the first medicine to be sold a a tablet. (4)

Acid-base titration

Aim(s)

� To find out how accurate acid-base titration (and the following calculations) are at determining the purity of a sample of aspirin of a known concentration.

� To find out which indicator would be best to use for this reaction

What is an acid-base titration?

An acid-base titration is an analytical method used to determine the concentration of an acid or base in aqueous solutions by neutralizing the acid/base in solution by a basic/acidic reagent. (5)

The concentration of a solution is how much of a chemical substance is dissolved in a given volume of solution. The equation for finding concentration is shown (left).

One solution is gradually added to another until the reaction between the two is complete. The point where the reaction ends is called the end point of the reaction and depends on the relative strengths of the acids and bases used. If a weak acid reacts with a weak base, a neutral solution can be obtained. However, any other combination will shift the endpoint either in the acidic direction or the basic direction. The completion of the reaction is usually shown by an indicator, added to the solution being titrated, through a colour change.

This method of analysis is appropriate for finding the concentration of aspirin because it is a weak acid and therefore can be neutralized using a base. In my experiment a known concentration of a strong base, sodium hydroxide of concentration 0.1moldm-3, is used. The reaction that takes place between the aspirin and the sodium hydroxide can be used to determine the concentration of aspirin in the sample, through the use of a simple calculation with the aid of the following reaction equation:

Acetylsalicylic Acid (Aspirin) + Sodium Hydroxide ? Sodium Acetylsalicylate + Water

(C9H8O4) (NaOH) (C9H7O4Na) ( H2O)

(3)

Risk assessment

Name of substance Quantity and/or concentration Risk Potential hazards Safety precautions taken

Sodium hydroxide (NaOH)

(7) 0.1moldm-3 Exothermic reaction with water-could eject hot solution.

Could be splashed into the eyes, onto the skin or clothes. Corrosive

Irritant Labratory coat

Labratory glasses

Acetylsalicylic acid (aspirin)

(7) Small amounts at one time-one tablet. Accumulation on fingers or powder inhaled. Harmful, especially if swallowed pure.

Irritating to eyes, respiratory system and skin. Labratory coat

Labratory glasses

Phenolphthalein indicator

(7)

Small amounts at one time, never coming into contact with skin. Can stain skin on contact. Irritant to eyes and respiratory system Labratory glasses

Labratory coat

Ethanol

(7) Small volumes at one time, never coming into contact with skin. Vapour will readily catch fire at temperatures above 13oC. Methanol is toxic by inhalation, if swallowed and by skin absorption so spills to be cleaned up quickly. Highly flammable

Contains methanol-toxic No naked flames

Labratory coat

Labratory glasses

Gloves (if in high concentration)

Preliminary work

Choosing an indicator

An acid-base indicator is a compound that changes colour depending on the pH of the solution it is placed in. There are many different acid-base indicators avaiLabratoryle and choosing which is appropriate for the reaction you are carrying out depends on a number of different things, such as:

� The relative strength of the acid and base used

This affects the position of equilibrium according to Le Chatiliers principle, which states that when a change is made to a system in equilibrium, the position of equilibrium will move in such a way to minimise the effects of the change. This is relevent to the equilibrium present in the indicator and the way in which a colour change is brought about.

� The equivilance point of the reaction

This is the point at which the number of moles of H+ ions is equal to the number of moles of OH- ions in a solution. In an acid-base titration this equivilance point will only be neutral if a strong acid and strong base are reacted together. However, in the acid-base titration of aspirin and sodium hydroxide, there is a weak acid and strong base reacting. This shifts the position of equilibrium to the basic side of the equation because the solution contains an aspirin salt which is basic. This means that the equivileance point will be above pH 7.0.

� The pH range on the indicator

Each indicator is capable of covering a specifc pH range. This means that depending on the relative strength of the acid and base used and the equivilance point of the reaction an indicator can be chosen which covers the required pH range. In the case of aspirin and sodium hydroxide, the equivilance point is above 7.0 and the indicators phenolphthalein and bromothymol blue are both active above this pH. Phenolphthalein, for example, has two forms. In acidic conditions, it is in the acid form, which is colourless. In basic conditions, a H+ ion is removed from each phenolphthalein molecule, converting it to its base form, which is pink.(9)

� The ease of detection of the colour change

For both phenolphthalein and bromothymol blue there is an easily detectable colour change:

Indicator Low pH High pH

Bromothymol blue yellow Blue

Phenolphthalein Colourless Pink

Chemicals and Equipment

Aspirin of known concentration (tablets)

Ethanol (95%)

Sodium hydroxide solution (0.1moldm-3)

Phenolphthalein indicator

Pestle and mortar

3 x Specimen bottle

2 x 250cm3 beaker

Access to electronic balance

3 x 100cm3 conical flask

10cm3 measuring cylinder

Burette (clamp and stand)

White tile

Glass funnel

Dropping pipette

Diagram

Method

1. Grind up one aspirin tablet using a pestle and mortar.

2. Transfer as much of the powder as possible to a specimen bottle, using a spatula. Weigh the specimen bottle to an accuracy of 0.01g and record the mass.

3. Using a measuring cylinder place 10cm3 of 95% ethanol into a 100cm3 conical flask. Add 5 drops of chosen indicator.

4. Transfer as much of the aspirin powder as possible from the specimen bottle to the conical flask. Re-weigh the specimen bottle and record the mass.

5. Swirl the conical flask until all the aspirin powder has dissolved. Do not let any of the solution splash out of the flask.

6. Titrate the solution in the flask with 0.1moldm-3 NaOH solution from a burette record the volume needed to produce the first tinge of pale pink colour in the indicator. This measures the end point of the titration.

7. Repeat the procedure until a minimum of three results within 0.01cm-3 is found. Find the average titre.

Final work

Chemicals and Equipment

Aspirin sample of known concentration (tablet form)

Sodium hydroxide (0.1moldm-3)

Ethanol (95%)

Phenolphthalein indicator

Bromothymol blue indicator

250cm3 volumetric flask

2 x 250cm3 beaker

6 x 250cm3 conical flask

100cm3 measuring cylinder

Specimen bottle

Volumetric pipette (and safety filler)

Burette (clamp and stand)

Volumetric flask bung

Distilled water

Access to an electronic balance

Glass funnel

Dropping pipettes

White tile

Pestle and mortar

Diagram

Making the aspirin solution

1. Using pestle and mortar crush one aspirin tablet and transfer as much of the powder as possible to the specimen bottle. Weigh specimen bottle and record mass.

2. Transfer as much of the powder as possible from the specimen bottle to a 250cm3 beaker. Reweigh specimen bottle and record mass. Find the mass of aspirin powder used (transferred to beaker)

3. Repeat steps 1 and 2 until the powder of 10 aspirin tablets has been added to the beaker. Find the total mass of aspirin powder used.

4. Add 100cm3 ethanol to the beaker containing the aspirin, using a 100cm3 measuring cylinder (measuring cylinder is accurate enough for this because the solution does not have to be exactly 50% ethanol, 50% distilled water, but approximately) and swirl to dissolve the aspirin powder. Using a clean measuring cylinder add 100cm3 distilled water to the beaker and swirl. Add another 25cm3 ethanol to the beaker and swirl.

5. Transfer this solution to a 250cm3 volumetric flask, with washings, using a clean funnel. Make up to the mark with distilled water. This should create a dissolved aspirin solution of 50% ethanol 50% distilled water.

Titration method

1. Rinse the burette with a small amount of 0.1moldm-3 NaOH solution. Fill the burette with the 0.1moldm-3 NaOH, overfill slightly so that some solution can run into a waste beaker filling the burette tip. Take the initial burette reading to the nearest 0.05cm3 and record this in a table.

2. Rinse a 25cm3 volumetric pipette with a small amount of your aspirin sample solution and run out into a waste beaker. Fill the volumetric pipette up to the mark, ensuring that the bottom of the meniscus rests exactly on the mark.

3. Carefully run the aspirin solution sample into a clean 250m3 conical flask. Once the pipette is empty, touch the tip against the inside of the conical flask. This ensures that the last drop is transferred from the pipette to the conical flask. The pipette has delivered exactly 25cm3 of the aspirin solution.

4. Add five drops of indicator (phenolphthalein or Bromothymol blue indicator � dependent on results from preliminary work), and swirl to mix.

5. Slowly run the NaOH from the burette into the conical flask swirling constantly and looking for the first hint of colour change given by the indicator. This first titration is called a rough titration so it does not matter if the end point is over shot. The purpose of the rough titration is to give an indication of the amount of NaOH. Record the final burette reading and find the volume of NaOH used (the titre).

6. If necessary, refill the burette and record the initial burette reading. It is best to begin on a whole number.

7. Using the volumetric pipette transfer 25cm3 of the aspirin solution to a clean conical flask. Add the same amount of indicator as the first time, 5 drops.

8. Run the NaOH slowly into the conical flask, swirling constantly, until you are within 5cm3 of the rough titre. Then add 0.5cm3 of NaOH to the conical flask, swirling after every addition until within 1cm3 of the rough titre.

9. Add the NaOH to the conical flask drop by drop, swirling after every addition, until a permanent change of colour is shown by the indicator. The colour should remain for 30s, if it does not continue to add NaOH to the conical flask until this occurs.

10. Record the end burette reading and calculate the titre. Repeat steps 6-9 until a minimum of 3 results within 0.10cm3 of each other is obtained. Find the average titre and record.

Back-titration

Aim

The aim of this experiment is to determine the percentage of aspirin in an aspirin tablet. The result will allow me to calculate the amount of aspirin present in each tablet and compare this to the amount of aspirin stated on the box. In completeing this analysis using a back titration method I will be able to compare the accuracy of back titration against acid-base titration and colorimetry.

What is a back titration?

Back titratioin undergoes a similar method to a forward titration, the difference is in that it is not carried out with the solution whose concentration is required to be known(the analyte) but with excess volume of reactant which has been left over after the completion of a reaction with the analyte.

When ingested, aspirin passes unchanged through the acidic conditions of the stomach but is hydrolysed in the alkaline juices of the intestines. This hydrolysis is mimacked in the labratory by hydrolysing aspirin with NaOH (1.0 moldm-3) and the equation for this reaction is shown below:

CH3COOC6H4 + NaOH ? CH3COONa + HOC6H4COONa + H2O

The unused sodium hydroxide which ramains after the hydrolysis can be titrated against a standard acid, in this experiment HCl (0.1moldm-3) is used. The amount of alkali required for the hydrolysis can now be calculated from the above equation, the number of moles of acetyl-salicylic acid which have been hydrolysed can be found.

Risk assessment

Risk assessment available on page 3 for most chemicals used in this experiment. Those that are not listed in the previous risk assessment can be found below:

Chemical Quantity and/or concentration Risk Possible hazards Safety precautions

Hydrochloric acid 1.0moldm-3

Small amounts, usually less than 50cm3 used at one time.

Could be splashed into the eyes, onto the skin or clothes.

Vapour could be inhaled if left uncovered. Irritant

Highly irritating to respiratory system Gloves

Labratory coat

Labratory glasses

Hydrochloric acid 0.1 moldm-3

Small volumes used with volumetric pipettes. No more than 50cm3 at one time. Could be splashed into the eyes, onto the skin or clothes.

Vapour could be inhaled if left uncovered. Irritant

Highly irritating to respiratory system Labratory coat

Labratory glasses

Sodium hydroxide

(NaOH) 0.1moldm-3

Small amounts, usually less than 50cm3 used at one time.

Exothermic reaction with water-could eject hot solution.

Could be splashed into the eyes, onto the skin or clothes. Corrosive

irritant Labratory coat

Labratory glasses

Gloves (if in high concentration)

Chemicals and Equipment

Aspirin of a known concentration (aspirin tablets)

NaOH (1.0moldm-3)

HCl (0.1moldm-3)

HCl (0.05moldm-3)

Phenolphthalein indicator

Anti-bumping granules

Volumetric pipette (and pipette filler)

Burette (clamp and stand)

2 x 250cm3 beaker

12 x 250cm3 conical flask

Distilled water

Funnel

White tile

100cm3 Pear shaped flask

Condenser

Bunsen burner

Tripod

Gauze

Heat resistant bench mat

Clamp stand and clamp

Access to running water

Spatula

Access to an electronic balance

Making hydrolysed aspirin solution

1. Using pestle and mortar crush one aspirin tablet and transfer as much of the powder as possible to the specimen bottle. Weigh specimen bottle and record mass.

2. Transfer as much of the powder as possible from the specimen bottle to a 100cm3 pear shaped flask. Reweigh specimen bottle and record mass. Find the mass of aspirin powder used (i.e. amount transferred to beaker).

3. Repeat steps 1 and 2 until a total mass of between 1.3 -1.7g has been added to the pear shaped flask.

4. Add 25cm3 ethanol to the pear-shaped flask, using a 25cm3 volumetric pipette and swirl to dissolve the aspirin powder. Using a clean 25cm3 volumetric pipette add 25cm3 1.0 NaOH to the pear shaped flask and swirl. Finally add 25cm3 distilled water to the pear shaped flask using a 25cm3 measuring cylinder.

5. Add a spatula of anti-bumping granules. (Do not stopper the flask)

Hydrolysis of the aspirin solution

Diagram

Method

1. Set up a Bunsen burner with tripod and gauze.

2. Attach a condenser vertically to a clamp stand, so that it �hangs� over the gauze and directly above the Bunsen burner.

3. Connect the condenser to a cold water supply with the water entering at the bottom and running out into a sink from the top.

4. Attach the pear-shaped flask to the bottom of the condenser in such a way that the bottom just rests on the gauze.

5. Using the Bunsen burner (on a blue-flame) heat the reactants so that they boil gently. The condensing vapour should ideally reach no more than half way up the condenser and the vapour should drip back into the pear-shaped flask no faster than 1 drop per second.

6. After 10-15 minutes turn off the Bunsen burner and allow the solution to cool. Once cool remove the pear-shaped flask from the condenser and carefully transfer the solution, with washing, to a 250cm3 volumetric flask taking care to leave the anti-bumping granules behind.

7. Make up to the mark with distilled water and stopped the flask.

Standardization titration

Diagram

Method

1. Rinse the burette with a small amount of 1.0moldm-3 NaOH solution. Fill the burette with the 1.0moldm-3 NaOH, overfill slightly so that some solution can run into a waste beaker filling the burette tip. Take the initial burette reading to the nearest 0.05cm3 and record this in a table.

2. Rinse a 25cm3 volumetric pipette with a small amount of 1.0moldm-3 HCl solution and run out into a waste beaker. Fill the volumetric pipette up to the mark, ensuring that the bottom of the meniscus rests exactly on the mark.

3. Carefully run the HCl solution sample into a clean 250m3 conical flask. Once the pipette is empty touch the tip against the inside of the conical flask, this ensures that the last drop is transferred from the pipette to the conical flask. Therefore the pipette has delivered exactly 25cm3 of the HCl solution.

4. Add five drops of phenolphthalein indicator and swirl to mix.

5. Slowly run the NaOH from the burette into the conical flask swirling constantly and looking for the first hint of colour change given (colourless to pink) by the indicator. This first titration is called a rough titration so it does not matter if the end point is over shot. The purpose of the rough titration is to give an indication of the amount of NaOH required. Record the final burette reading and find the volume of NaOH used (the titre).

6. If necessary, re-fill the burette and record the initial burette reading. It is best to begin on a whole number.

7. Using the volumetric pipette transfer 25cm3 of HCl solution to a clean conical flask. Add the same amount of indicator as the first time, 5 drops.

8. Run the NaOH slowly into the conical flask, swirling constantly, until you are within 5cm3 of the rough titre. Then add 0.5cm3 of NaOH to the conical flask, swirling after every addition until within 1cm3 of the rough titre.

9. Add the NaOH to the conical flask drop by drop, swirling after every addition, until a permanent change of colour is shown by the indicator. The colour should remain for 30s, if it does not continue to add NaOH to the conical flask until this occurs.

10. Record the end burette reading and calculate the titre. Repeat steps 6-9 until a minimum of 3 results within 0.10cm3 of each other is obtained. Find the average titre and record.

Hydrolysis of aspirin titration

Diagram

Method

1. Rinse the burette with a small amount of 0.1moldm-3 HCl solution. Fill the burette with the 0.1moldm-3 HCl, overfill slightly so that some solution can run into a waste beaker filling the burette tip. Take the initial burette reading to the nearest 0.05cm3 and record this in a table.

2. Rinse a 25cm3 volumetric pipette with a small amount of your hydrolysed aspirin solution and run out into a waste beaker. Fill the volumetric pipette up to the mark, ensuring that the bottom of the meniscus rests exactly on the mark.

3. Carefully run the hydrolysed aspirin solution sample into a clean 250m3 conical flask. Once the pipette is empty, touch the tip against the inside of the conical flask, this ensures that the last drop is transferred from the pipette to the conical flask. Therefore the pipette has delivered exactly 25cm3 of the hydrolysed aspirin solution.

4. Add five drops of phenolphthalein indicator, and swirl to mix. The solution should turn a pink colour.

5. Slowly run the HCl from the burette into the conical flask swirling constantly and looking for the first hint of the indicator colour fading. This first titration is called a rough titration so it does not matter if the end point is over shot. The purpose of the rough titration is to give an indication of the amount of HCl. Record the final burette reading and find the volume of HCl used (the titre).

6. If necessary, refill the burette and record the initial burette reading. It is best to begin on a whole number.

7. Using the volumetric pipette transfer 25cm3 of the aspirin solution to a clean conical flask. Add the same amount of indicator as the first time.

8. Run the HCl slowly into the conical flask, swirling constantly, until you are within 5cm3 of the rough titre. Then add 0.5cm3 of HCl to the conical flask, swirling after ever addition until within 1cm3 of the rough titre.

9. Add the HCl to the conical flask drop by drop, swirling after every addition, until a there is a permanent change; pink to colourless shown by the indicator. The colour should not return for 30s, if the colour returns within this time continue to add HCl to the conical flask until it is colourless.

10. Record the end burette reading and calculate the titre. Repeat steps 6-9 until a minimum of 3 results within 0.10cm3 of each other is obtained. Find the average titre and record.