Task
The objective of this experiment is to determine the % of
2 – Ethanoylhydroxybenzoic acid in aspirin tablets.
A known amount of NaOH (Sodium Hydroxide) solution is used to hydrolyse a known mass of aspirin tablets.
CH3COOC6H4COOH + 2NaOH → CH3COONa + HOC6H4COONa + H20
Standard acid (0.1-mol dm-3 hydrochloric acid) is titrated with the remains of the unused sodium hydroxide. So, the amount of alkali needed for the hydrolysis can be calculated from the equation. Also, the number of moles of the aspirin acid (which have hydrolysed) can be found.
Equipment
- Aspirin tablets (about 5)
- 25 cm³ pipette and filler
- Graduated flask (250 cm³)
- Bunsen burner, tripod, heat-proof mat, gauze
-
1.0-mol dm-3 NaOH (about 30 cm³)
-
0.1-mol dm-3 hydrochloric acid (about 150 cm³)
- Phenolphthalein indicator
Method
Part 1(with diagram) – Measuring the amount of acid that reacts with 25 cm³ of unreacted alkali
-
With the aid of a safety filler, using a 25-cm³ pipette, obtain exactly 25 cm³ of the NaOH (1-mol dm-3) solution.
- Put this into a 250 cm³-graduated flask.
-
Go up-to the mark on the flask with distilled water (H2O).
- Obtain 25 cm³ of this solution (using a pipette) and drop into a beaker. Add a couple of drops of Phenolphthalein indicator (solution should go purple)
-
Place this beaker below the Burette – which should be filled with 0.1-mol dm-3 HCl acid.
- Carry out the titration until the solution goes colourless (keep shaking the beaker whilst releasing the HCl acid).
Part 2 – Hydrolysis of Aspirin
- Weigh out accurately 5 aspirin tablets and put them into a conical flask.
-
Drop 25cm³ of NaOH (1.0-mol dm-3) onto the tablets (with aid of a pipette)
- To hydrolyse the aspirin acid – heat it gently for 10 minutes.
SAFETY: eye protection must be worn.
- Cool the mixture and transfer (with washings) to a 250 cm³-graduated flask. Fill up to the blue mark with distilled water.
Part 3 – Estimating the quantity of unused NaOH after hydrolysis
- Put 25 cm³ of the hydrolysed solution (from part 2) into a conical flask using a pipette.
- Put a couple of drops of Phenolphthalein indicator in the solution. It should go purple.
-
Titrate this against 0.1-mol dm-3 HCl acid (which is in the Burette).
- Carry out the titration until the solution goes colourless (keep shaking the beaker whilst releasing the HCl acid).
Results
Part 1:
Average (t1) =
(25.25 + 25.30)
2
= 25.275 cm³ of HCl acid
Part 2:
Mass of aspirin tablets = 1.66 g
Part 3:
Average (t2) =
(7.80 + 7.65)
2
= 7.725cm³ of HCl acid
The amount of Aspirin per tablet
The volume of acid required in part 1 (t1) = 25.275 cm³
The volume of acid required in part 1 (t2) = 7.725 cm³
Number of moles of NaOH that react with aspirin:
t1 – t2
1000
= 25.275 – 7.725
1000
= 0.01755 moles (NaOH)
The equation say there are half as many moles of aspirin as NaOH. Therefore, moles of pure aspirin:
t1 – t2
2000
= 25.275 – 7.725
2000
= 0.008775 moles
Relative Molecular Mass (RMM) of Aspirin:
CH3COOC6H4COOH
Mr = 180
Mass of pure aspirin (Y grams):
0.008775 x 180 = 1.5795 g
The % purity of the tablets:
Y x 100
Mass of tablets
= 1.5795 x 100
1.66
=95.15060241
= 95.15% (2d.p.)
Amount of aspirin per tablet
Mass per tablet:
1.66
5
= 0.332 g
Mass of pure aspirin per tablet (Z):
1.5795
5
= 0.3159
The % purity per tablet:
Z x 100
Mass of tablets
= 0.3159 x 100
0.332
=98.10559006…
= 98.11% (2d.p.)
Evaluation
I feel that the experiment was a general success. The techniques used in the experiment were all useful and necessary. They did not raise any problems whilst the experiments were being carried out.
So, I feel there does not need to be any changes to the techniques if the experiment was repeated.
There were not any major errors whilst the experiment was being carried out, so I can say with confidence that the results are reliable. The tables in which the results were recorded did not cause any problems.
However, minor errors may have occurred in these areas:
Quantifiable Errors:
Balance: There could have been a slight error whilst the aspirin tablets were being weighed. This is due to the fact that balances can never be accurate. An example is a balance being accurate to 0.001g. This shows the balance must round off to the nearest 1/100th of a gram, which will cause a slight error.
Burette: This is accurate to 0.05 cm³. If the reading has to be rounded because the scale on the Burette is not big enough to be exact, this will also cause an error.
Graduated Flask: If the volume of the flask is 250 cm³, the bottom of the meniscus is on the calibration line, the error is 0.1 cm³.
Pipette: Once again, the scale is not extremely accurate and may cause an error.
Procedural errors:
It could have been difficult to spot the endpoint of the reaction. Despite the solution being purple, and the colour looked for was colourless – it may be hard to get the exact amount of HCl added to go colourless. This is because the releaser of the Burette may release too much of the HCl acid when you know you only need a small amount of it for the solution to go colourless.
The equipment may not have washed well (drying is also very important).
Some of the HCl acid may have got onto the side of the conical flask and not reacted with the hydrolysed aspirin.
If an acid-base titration had been carried out, it would not have been a good idea to put too much indicator in, because the indicator itself is a weak acid.
Bibliography
- Times Education As Chemistry CD ROM
- My notes
- Internet
- Class text book
