Phenolphthalein indicator was used in the following reaction, because it is a suitable indicator for a titration between a weak acid and a strong alkali.
Requirements
- Aspirin
-
100 cm3 conical flask
- 95% ethanol
-
Pipette (10 cm3)
-
Sodium hydroxide solution 0.1 mol dm-3
- Burette
- Phenolphthalein indicator
- Analytical balance (4 d.p)
- Burette
- Burette Stand
-
Volumetric flask 500 cm3
- Distilled Water
Funnel
White Tile
Goggles
Laboratory coat
Apparatus Justification
- Aspirin - contains the acetylsalicylic acid, which will react with the sodium hydroxide
- Burette – This is used to add the sodium hydroxide. It makes the results as accurate as possible, as it allows the sodium hydroxide to be added drop by drop when the solution is close to neutralising.
- Burette Stand – This holds the burette steady in place.
- Conical Flask – This is used to mix the aspirin with the ethanol. The shape of it makes it less likely that any solution should spill out.
-
Phenolphthalein indicator- The colour of this indicates when the endpoint has been reached.
-
Sodium Hydroxide 0.1 mol dm-3 – This will hydrolyse the aspirin into salicylic and acetate ions
-
Volumetric flask 500 cm3 – This is used to mix up solutions accurately.
- Distilled Water – This is used to make up the correct solutions and to wash equipment. Normal water will not do because of the impurities in it.
- Funnel – This allows hydrochloric acid to be poured into the burette so none is spilt.
- Pipette 10cm³ – This is used to accurately draw out the correct amount of a substance from a solution.
- White Tile – So the colour of the solution is easier to see.
- Goggles – Used to protect the eyes throughout the experiment.
- Laboratory coat – Used to protect the skin and clothing.
- Analytical balance- used to way the aspirin and sodium hydroxide accurately to 4 decimal places.
Safety precautions (9)
Results obtained were recorded in a table format:-
Recrystalised aspirin
Mass of aspirin used
Accurate 1 = -------g
Accurate 2 = -------g
Accurate 3 = -------g
Pure aspirin
Mass of aspirin used
Accurate 1 = --------g
Accurate 2 = --------g
Accurate 3 = --------g
BACK TITRATION (10) (14) (15) (16)
Introduction
Aspirin is an analgesic drug that can be harmful if taken in excess. To find out about the purity of aspirin, a technique called back titration is used. Aspirin is a weak acid that undergoes slow hydrolysis. Therefore to overcome the problem, a known, mass of Aspirin is hydrolyzed with a known excess amount of base i.e. in this case NaOH. The excess is then back-titrated with a standard solution of HCl. The number of moles of NaOH required for the hydrolysis can then be calculated, and the number of moles of acetyl Salicylic acid that has been hydrolyzed can also be calculated.
Requirements
- Aspirin
- Burette
- Burette stand
- Conical flask
- Methyl red indicator
-
Hydrochloric acid 0.1 mol dm-3
-
Sodium hydroxide 1.0 mol dm-3
-
Volumetric flask (250 cm3)
- Distilled water
- Funnel
-
Pipette (25 cm3)
- Analytical balance
- White Tile
- Goggles
- Laboratory coat
- Burette Stand
- Analytical balance ( 4 d.p)
Apparatus Justification
- Aspirin - contains the acetylsalicylic acid, which will react with the sodium hydroxide
- Burette – This is used to add the hydrochloric acid to the sodium hydroxide. It makes the results as accurate as possible, as it allows the hydrochloric acid to be added drop by drop when the sodium hydroxide is close to neutralising.
- Burette Stand – This holds the burette steady in place.
- Conical Flask – This is used to react the aspirin with the sodium hydroxide. The shape of it makes it less likely that any solution should spill out.
- Methyl red indicator – The colour of this indicates when the sodium hydroxide has been neutralised by the hydrochloric acid.
-
Hydrochloric Acid 0.1 mol dm-3 – This will neutralise the sodium hydroxide
-
Sodium Hydroxide1.0 mol dm-3 – This will hydrolyse the aspirin into salicylic and acetate ions
-
Volumetric flask 250 cm3 – This is used to mix up solutions accurately.
- Distilled Water – This is used to make up the correct solutions and to wash equipment. Normal water will not do because of the impurities in it.
- Funnel – This allows hydrochloric acid to be poured into the burette so none is spilt.
- Pipette 25cm³ – This is used to accurately draw out the correct amount of a substance from a solution.
- White Tile – So the colour of the solution is easier to see.
- Goggles – Used to protect the eyes throughout the experiment.
- Laboratory coat – Used to protect the skin and clothing.
- Analytical balance- used to way the aspirin and sodium hydroxide accurately to 4 decimal places.
Safety Precautions (9)
Results obtained were recorded in a table format:-
Recrystalised aspirin
Mass of aspirin used = -------g
Weight of container = --------g
Pure aspirin
Mass of aspirin used = -------g
Weight of container = --------g
Synoptic grid
Preparation of Aspirin
Apparatus
- Salicylic acid
- Ethanoic anhydride
- Concentrated Sulphuric acid
- Ethanoic acid (glacial)
-
100 cm3 conical flask
- Distilled Water
-
10 cm3 pipette
- 100cm3 beaker
- Hirsch funnel
- Water bath containing crushed ice
- Goggles
- Laboratory coat
- Analytical balance ( 4 d.p)
Method
-
4.0029g of 2-hydroxbenzoic acid was weighed out on a 4.d.p balance and placed in a 100 cm3 conical flask. (4.0029g was used to ensure that enough aspirin was made so that further experiments could be carried out.)
-
8cm3 of ethanoic anhydride was added to the conical flask using a 10cm3 pipette making sure that the bottom of the meniscus is on the graduation mark (this will hydrolyse the ethanoic acid).
- 10 drops of concentrated sulphuric acid was added to the flask (this acts as a catalyst). The flask was swirled for about 15 minutes making sure that the contents were mixed thoroughly. Crystals of Aspirin started appearing and a crystalline mush was formed.
-
The crystalline mush was diluted by adding 8cm3 of cold glacial ethanoic acid using a 10cm3 pipette and cooled by placing in a water bath containing crushed ice (product becomes less soluble when its cool). This was left for approximately 10 – 15 minutes until an impure solid was formed.
- The crystals were then filtered off using a Hirsch funnel (a small funnel for vacuum filtration) washing a few times with ice cold water (this removes some ethanoic acid as it is water soluble). All the Aspirin in the beaker was washed a few times and poured into the Hirsch funnel so that no crystals were lost.
-
The crude aspirin was placed in a 100cm3 beaker and hot water was added to dissolve the aspirin. The crystals were then cooled and filtered using a Hirsch funnel, all the Aspirin in the beaker was washed a few times with cold water (hot water is not used because the crystals would dissolve) and poured into the Hirsch funnel so that no crystals were lost this process is known as re-crystallisation and is a way of purifying a solid product.
- The crystals were then placed on a filter paper to dry overnight.
- Once the crystals were dry, they were weighed and the theoretical yield and percentage yield of aspirin was calculated.
Results
Appearance of crystals = white crystals
Amount of salicylic acid used = 4.0029g
Amount of re-crystallised product produced = 2.6586g
Determine the melting point of Aspirin
Apparatus
- Melting point apparatus
- Capillary tubes
- Thermometer
- Aspirin
- Mortar and pestle
Method
- Melting point apparatus was used to determine the melting point of the re-crystallised aspirin.
- Using a mortar and pestle, the re-crystallised aspirin was pulverized into a fine pile in the mortar.
-
The open end of a capillary tube was pushed into the pile of aspirin powder. The aspirin was packed into the capillary tube to a depth of about 1cm3 by tapping lightly on the table top.
- A thermometer was placed in the melting point apparatus and the temperature was set to approximately 40V.
- The capillary tube was placed in the melting point apparatus and the crystals were observed through a magnifying glass of the apparatus.
- The range at which the aspirin melted was noted.
The experiment was then repeated for pure aspirin.
Results
Forward titration
Apparatus
- Aspirin
-
100 cm3 conical flask
- 95% ethanol
-
Pipette (10 cm3)
-
Sodium hydroxide solution 0.1 mol dm-3
- Burette
- Phenolphthalein indicator
- Analytical balance (4 d.p)
- Burette
- Burette Stand
-
Volumetric flask 500 cm3
- Distilled Water
Funnel
White Tile
Goggles
Laboratory coat
Preparation of 0.1 mol dm-3 sodium hydroxide solution
A standard solution is one whose concentration is known exactly. Standard solutions can be prepared by weighing a mass of solid, and dissolving it with known volume of solution in a standard flask. Standard solutions can be chemically reacted with a solution of unknown concentration in order to determine the concentration of the unknown
The acetylsalicylic acid and NaOH react in a 1:1 ratio.
The NaOH solution was made to a concentration of 500cm3 0.1 mol dm-3. The next step was to calculate how much sodium hydroxide was needed to make a 0.1moldm-3 solution:-
First the relative atomic mass of sodium hydroxide was calculated:
(Relative Atomic Masses: Na = 23, O = 16)
Mr (NaOH) = 23 + 16 + 1 = 40g
1 mol dm of NaOH = 40g
0.1 mol dm of NaOH = 40
10
= 4g
To make 1000cm3 of a 0.100 mol dm-3 solution of sodium hydroxide you would need 4g of the solid. So therefore to make 500cm3 of a 0.1 mol dm-3 solution of sodium carbonate, 2.00g was needed.
Into a weighing bottle 2.00g of solid sodium hydroxide was weighed accurately and the mass was then recorded. An analytical balance was used to weigh the solid because they are very accurate. A clean beaker was rinsed with distilled water and the solid was transferred from the weighing bottle into the beaker. The weighing bottle was rinsed three times with distilled water, transferring the washings each time. This was done to ensure that the entire solid goes into the beaker. Approximately 100 cm3 of distilled water was added to the beaker. This was then stirred using a glass rod to dissolve the solid. The solution was then washed of the glass rod to ensure none is lost. The 500cm3 standard flask was washed three times and using a filter funnel the solution was carefully transferred in to it. The beaker and funnel was then rinsed several times making sure that all the solution has gone into the standard flask. Distilled water was then added until the solution was 1cm3 below the graduation mark. Using a dropping pipette enough water was added to bring the bottom of the meniscus up to the graduation mark. A stopper was added to the flask and inverted several times to help mix the contents.
Method
Once the solution was prepared the next procedure conducted was the titration. In this method measurements were taken carefully in order to accurately obtain precise results in the titration.
Preparing the burette
The burette was washed with distilled water; checks were made so that there were no leaks. Using a small funnel, 5-10cm3 of the prepared solution of 0.1mol dm-3 sodium hydroxide was added to the burette. The funnel was then removed. The burette was then taken from the stand and tipped and rotated to wash the inside surface with the solution. The solution was then placed into the waste beaker.
The burette was filled using a funnel with the prepared solution of sodium hydroxide so the meniscus was above the zero mark. Accurate reading was taken by placing a white piece of paper behind the scale. The waste beaker was placed underneath the burette and the tap was opened until the solution filled the jet, it was made sure no air bubbles were present.
The solution was then allowed slowly to run until the bottom of the line of the meniscus was on the zero mark.
Using a clean dry weighing bottle, three samples of the re-crystalised aspirin was weighed to an approximate 0.3000g.
Using a pipette 10cm3 of 95% ethanol was added to a 100cm3 conical flask followed by 3 drops of phenolphthalein indicator. Aspirin was then added to the flask (as much as possible) and swirled gently to dissolve it. Care was taken so that no solution splashed out of the flask. (Aspirin is not very soluble in water therefore ethanol is used to help it dissolve)
Performing a rough titration
A rough titration was carried in order to distinguish an approximate endpoint
In order to perform a rough titration a note of the burette reading was recorded to the nearest 0.05cm3. The conical flask was placed on a white tile under the burette (white tile was used to view the end point clearly) Then from the burette the 0.1 mol dm-3 sodium hydroxide solution was added 1cm3 at a time, until there was a colour change from colourless to the first tinge of pale pink. This indicated the end point of the reaction. A note of the new burette reading was then recorded. To work out the volume of the solution added a calculation was made whereby the initial reading was subtracted from this end point value. The value calculated was the titre. The first titre was only a guide; therefore it was not included when the average value was calculated.
Performing an accurate titration
In order to perform an accurate titration, the whole titration method was carried out again. As the rough end point was approached, the 0.1 mol dm-3 sodium hydroxide solution was added a drop at a time and the flask was shaken. Each drop was added until one drop caused the colour to change. The titration was repeated to get concordant results within 0.1cm3 of each other. Results obtained were recorded in a table format.
The whole experiment was repeated again, but this time using pure aspirin.
Results
Re-crystalised aspirin
Mass of aspirin used
Accurate 1 = 0.3050g
Accurate 2 = 0.3038g
Accurate 3 = 0.3007g
Pure aspirin
Mass of aspirin used
Accurate 1 = 0.3178g
Accurate 2 = 0.3047g
Accurate 3 = 0.3032g
BACK TITRATION
Apparatus
- Aspirin
- Burette
- Burette stand
- Conical flask
- Methyl red indicator
-
Hydrochloric acid 0.1 mol dm-3
-
Sodium hydroxide 1.0 mol dm-3
-
Volumetric flask (250 cm3)
- Distilled water
- Funnel
-
Pipette (25 cm3)
- Analytical balance
- White Tile
- Goggles
- Laboratory coat
- Burette Stand
- Analytical balance ( 4 d.p)
Preparation of 1.0 mol dm-3 sodium hydroxide solution
A standard solution is one whose concentration is known exactly. Standard solutions can be prepared by weighing a mass of solid, and dissolving it with known volume of solution in a standard flask. Standard solutions can be chemically reacted with a solution of unknown concentration in order to determine the concentration of the unknown
The acetylsalicylic acid and NaOH react in a 1:1 ratio.
The NaOH solution was made to a concentration of 250cm3 1.0 mol dm-3. The next step was to calculate how much sodium hydroxide was needed to make a 1.0 mol dm-3 solution:-
First the relative atomic mass of sodium hydroxide was calculated:
(Relative Atomic Masses: Na = 23, O = 16)
Mr (NaOH) = 23 + 16 + 1 = 40g
1 mol dm of NaOH = 40g
1.0 mol dm of NaOH = 40
4
= 10g
To make 1000cm3 of a 1.0 mol dm-3 solution of sodium hydroxide you would need 40g of the solid. So therefore to make 250cm3 of a 1.0 mol dm-3 solution of sodium hydroxide, 10g was needed.
Into a weighing bottle 10.00g of solid sodium hydroxide was weighed accurately and the mass was then recorded. An analytical balance was used to weigh the solid because they are very accurate. A clean beaker was rinsed with distilled water and the solid was transferred from the weighing bottle into the beaker. The weighing bottle was rinsed three times with distilled water, transferring the washings each time. This was done to ensure that the entire solid goes into the beaker. Approximately 100cm3 of distilled water was added to the beaker. This was then stirred using a glass rod to dissolve the solid. The solution was then washed of the glass rod to ensure none is lost. The 250cm3 standard flask was washed three times and using a filter funnel the solution was carefully transferred in to it. The beaker and funnel was then rinsed several times making sure that all the solution has gone into the standard flask. Distilled water was then added until the solution was 1cm3 below the graduation mark. Using a dropping pipette enough water was added to bring the bottom of the meniscus up to the graduation mark. A stopper was added to the flask and inverted several times to help mix the contents.
Method
Once the solution was prepared the next procedure conducted was the titration. In this method measurements were taken carefully in order to accurately obtain precise results in the titration.
Preparing the burette
The burette was filled with distilled water; checks were made so that there were no leaks. Using a small funnel, 5-10cm3 of the prepared solution was added to the burette. The funnel was then removed. The burette was then taken from the stand and tipped and rotated to wash the inside surface with the solution. The solution was then placed into the waste beaker.
The burette was filled using a funnel with the prepared solution so the meniscus was above the zero mark. Accurate reading was taken by placing a white piece of paper behind the scale. The waste beaker was placed underneath the burette and the tap was opened until the solution filled the jet, it was made sure no air bubbles were present. The solution was then allowed slowly to run until the bottom of the line of the meniscus was on the zero mark.
- Approximately 1.5g of aspirin was accurately weighed out in a conical flask using an analytical balance.
-
25 cm3 of 1.0 mol dm-3 sodium hydroxide solution was added to the conical flask and aspirin.
-
25 cm3 of distilled water was also added to the conical flask.
- The contents were simmered gently for 10 minutes. This was done by hydrolyzing the aspirin( acetyl salicylic acid)
-
The solution was then transferred into a 250 cm3 volumetric flask. The conical flask was rinsed a few times to ensure that all the aspirin has been transferred to the volumetric flask. The solution in the volumetric flask was then made up to 250 cm3 using distilled water.
-
25cm3 of the hydrolyzed solution was then pipetted into a conical flask.
-
4 drops of methyl red indicator was added to the 25cm3 hydrolyzed solution.
-
The solution was then titrated with 0.1 mol dm-3 hydrochloric acid.
Performing a rough titration
A rough titration was carried in order to distinguish an approximate endpoint
In order to perform a rough titration a note of the burette reading was recorded to the nearest 0.05cm3. The conical flask was placed on a white tile under the burette (white tile was used to view the end point clearly) Then from the burette the solution was added 1cm3 at a time, until there was a colour change from pink to colourless. This indicated the end point of the reaction. A note of the new burette reading was then recorded. To work out the volume of the solution added a calculation was made whereby the initial reading was subtracted from this end point value. The value calculated was the titre. The first titre was only a guide; therefore it was not included when the average value was calculated.
Performing an accurate titration
In order to perform an accurate titration, the whole titration method was carried out again. As the rough end point was approached, the solution was added a drop at a time and the flask was shaken. Each drop was added until one drop caused the colour to change. The titration was repeated to get concordant results within 0.1cm3 of each other. Results obtained were recorded in a table format.
The titrations were repeated until there were concordant results( 0.1 cm3 of each other).
The experiment was then repeated for pure aspirin.
Results
Re-crystalised aspirin
Mass of aspirin used = 1.5161g
Weight of container = 0.0000g
Pure aspirin
Mass of aspirin used = 1.5067g
Weight of container = 0.0000g
Preparation of Aspirin (12)
Calculating the percentage yield
% yield = actual yield x 100
Theoretical yield
Mr of salicylic acid = C7H603
12 x 7 + 1 x 6 + 16 x 3 = 138
Mr of Aspirin (acetylsalicylic acid) = C9H8O4
12 x 9 + 1 x 8 + 16 x 4 = 180
4.0029g of salicylic acid was used. Therefore the amount of acetylsalicylic acid produced should have been
Number of moles = mass
Mr
= 4.0029
138
= 0.0290 mol dm
According to the equation there is a 1:1 ratio which means 0.0290 moles of salicylic acid should produce 0.0290 moles of aspirin.
Therefore the theoretical yield of aspirin:
0.0290 x 180
= 5.22g
But actual yield was 2.6586g
Therefore the percentage yield = actual yield
Theoretical yield
= 2.6586 x 100
5.22
= 50.93%
Results
Forward titration(12)
Calculations
Re-crysatalised aspirin
Average mass off aspirin = 0.3032g
Average volume = 16.3000cm3
Number of moles of NaOH
N = c x v
= 0.1 x 16.30
1000
= 0.00163 mol dm
The mass that reacted with aspirin:
M = n x Mr
= 0.00163 x 180
= 0.2934g
The mass of aspirin that didn’t react was:
0.3032 – 0.2934
= 0.0098g
Percentage by mass of aspirin (percent purity) = 0.2934 x 100 = 96.76%
0.3032
Pure aspirin
Average mass off aspirin = 0.3086g
Average volume = 16.7250cm3
Number of moles of NaOH
N = c x v
= 0.1 x 16.7250
1000
= 0.0016mol dm
The mass that reacted with aspirin:
M = n x Mr
= 0.0016 x 180
= 0.3011g
The mass of aspirin that didn’t react was:
0.3086 – 0.3011
= 0.0075g
Percentage by mass of aspirin (percent purity) = 0.3011 x 100 = 97.56%
0.3086
BACK TITRATION(12)
Calculations
Re-crysatalised aspirin
Average titre = 8.50 + 8.50 = 8.50cm3
2
Mass of aspirin used = 1.5067g
Mass of container = 0.0000g
Number of moles present in 0.1 mol dm-3 HCL
N = c x v
= 0.1 x 8.50
1000
= 0.00085 moles (present in 25cm3 of NaOH)
Number of moles present in 1.0 mol dm-3 NaOH
N = c x v
= 1.0 x 25
1000
= 0.0250 moles
The number of moles that reacted with aspirin
0.0250 – 0.0085 (present in 250cm3 of NaOH)
= 0.0165 moles
The reaction is 1:2 ratio. One mole of aspirin needs two moles of NaOH.
Therefore the number of moles of aspirin used:
0.0165 = 0.0083 moles
2
The mass of aspirin that reacted with NaOH
M = n x Mr
= 0.0083 x 180
= 1.4850g
Therefore the mass of aspirin that didn’t react with NaOH.
1.5067 – 1.4850
= 0.0217g
Percentage by mass of aspirin (percent purity) = 1.4850 x 100 = 98.56%
1.5067
Pure aspirin
Average titre = 8.45 x 8.40 = 8.4250cm3
2
Mass of aspirin used = 1.5067g
Mass of container = 0.0000g
Number of moles present in 0.1 mol dm-3 HCL
N = c x v
= 0.1 x 8.4250
1000
= 0.00084moles (present in 25cm3of NaOH)
Number of moles present in 1.0 mol dm-3 NaOH
N = c x v
= 1.0 x 25
1000
= 0.0250 moles
The number of moles that reacted with aspirin
0.0250 – 0.0084(present in 250cm3 of NaOH)
= 0.0166 moles
The reaction is 1:2 ratio. One mole of aspirin needs two moles of NaOH.
Therefore the number of moles of aspirin used:
0.0166 = 0.0083 moles
2
The mass of aspirin that reacted with NaOH
M = n x Mr
= 0.0083 x 180
= 1.4850g
Therefore the mass of aspirin that didn’t react with NaOH.
1.5067 – 1.4850
= 0.0217g
Percentage by mass of aspirin (percent purity) = 1.4850 x 100 = 98.56%
1.5067
Conclusion
The aim of the investigation was to prepare aspirin and determine its purity by using different experimental techniques. The purity of the prepared aspirin was then compared to the purity of pure aspirin via the same experimental techniques.
The first stage was to synthesise aspirin (acetylsalicylic acid) from salicylic acid. There is a 1:1 ration between salicylic acid and acetyl salicylic acid. This means that the amount of salicylic acid used should be equal to the amount of acetylsalicylic produced. The results show that the theoretical yield should be 5.22g but the actual yield was 2.6586g. This gave a percentage yield 50.93%.
The experiment did go to plan, because white crystals were produced, and aspirin in its pure form is a white crystalline solid. The yield was lower than expected which could have been errors such as loss of crystals during filtration and washing. Also the salicylic acid may not have a fully dissolved in the ethanoic anhydride. This would affect the overall yield of aspirin because not all of the salicylic acid was synthesized.
Further experiments were then carried out to determine how pure the product was. The melting point of the re-crystallised aspirin was compared to the melting point of pure aspirin.
The melting point of the re-0crystallised aspirin was 130-135oC, which was lower than that of pure aspirin whose melting point was 142-142oC.
The low melting point of the re-crystallised aspirin suggests that there were impurities present in the product. However the melting point range of the re-crystallised aspirin was not far of the melting point of pure aspirin. This suggests that the concentration of impurities present was very low. The most likely sources of impurities could be the presence of water in the sample as it is very difficult to remove water from a product completely.
The best way to determine the amount of acid salicylic acid present was to do titrimetric analysis. The first titration to be carried out was forward titration. This is an analytical method used to determine the strength of a solution, or the concentration of a substance in solution. It allows you to determine the amount of reactant in the titration flask.
This is an alkaline hydrolysis known as sponification. It is the reaction of an ester with sodium hydroxide to produce an alcohol and the sodium salt of the carboxylic acid of the ester.
The purity of the re-crystallised aspirin was 96.76% compared with the purity of pure aspirin which was 97.56%. This suggests that the re-crystallised aspirin is pre with very little impurities.
However if impurities are present, forward titration will neutralise acetylsalicylic acid and the impurities as well. Also using this titration it was difficult to identify the end point because aspirin is a weak acid and the reaction takes place very slowly.
The final method used was back titration. This is an analytical method used, in which a known mass of aspirin is hydrolysed with excess amount of base, in this case NaOH. The excess is then back titrated with a standard solution of HCL.
The purity of the re-crystallised aspirin and pure aspirin was 98.56%. This suggests that the re-crystallised aspirin is pure.
After conducting all the experiments, it can be conducted that the back titration method was the most accurate. The reason is because;
- The endpoint is more easily recognized because the reaction is between a stronger base and a strong acid.
- The reaction is faster
- It gave reliable results
- The number of moles present in aspirin was easily calculated and the highest percentage yield that reacted was achieved.
Evaluating results and procedures
Introduction
The evaluation is a critical look at the method, results and conclusions and deciding how accurate and how reliable they are. This section also Discusses how the procedure could have been improved to give a more accurate and reliable conclusion.
There are many types of error that can occur during a practical investigation such as procedural error and precision error.
The accuracy normally depends upon the type of equipment being used, and how easy it was to take the measurements. The accuracy can vary throughout the investigation and this can have an effect over the range of results taken.
Accuracy of Results
From my results it can be concluded that there were no anomalous results and that the results were concordant. An anomalous result is something that you don't expect; an irregularity. If the results were not 0.1cm of each other then they could be classified as anomalies. Anomalies can occur from simply not measuring things properly. Although there were no anomalous results, this does not indicate that there were no errors.
Reliability and precision
There is a source of uncertainty in the precision of the apparatus which was used. Each measuring instrument is designed to measure to a certain level of precision. To compare the importance of the precision errors for different measuring instruments, a percentage error calculation was carried out using the following formulae:-
Percentage error = Error x 100
Reading
Balance error
An analytical balance was used to weigh the mass of solid
+ 0.00005g is the precision error of the balance.
Volumetric flask (class B)
If a 250 cm3 volumetric flask is filled correctly, so that the bottom of the meniscus rests on the calibration line, the error is 0.2cm3.
Volumetric flask (class B)
If a 500 cm3 volumetric flask is filled correctly, so that the bottom of the meniscus rests on the calibration line, the error is 0.25cm3.
Burette (class B)
All burettes read to 2 decimal places in which the second figure is either 0 or 5. One drop from a burette has a volume approximately 0.05 cm3. Therefore the precision error is 0.05.
Pipette (class B)
The error on a 25cm3 pipette is 0.06 cm3 if used correctly, i.e. it is allowed to drain and retain its last drop
.
Pipette (class B)
The error on a 10cm3 pipette is 0.02 cm3 if used correctly, i.e. it is allowed to drain and retain its last drop.
Melting point apparatus
The thermometer creates an error of 0.1
Measuring cylinders and beakers were not used during this investigation due to greater uncertainty associated with them.
The total percentage error can then be calculated by adding all the errors up
.
Synthesis of aspirin
Balance error
Percentage error = 0.00005 x 100 = + 0.00125%
4.0029
Pipette (class B) 10 cm3
Percentage error = 0.10 x 100 = + 1.25%
8
Total percentage error = 0.00125 + 1.25
= + 1.25125%
By working out the total percentage error, the confidence limit can be calculated.
Actual weight of salicylic acid = 4.0029g
Total % error = 1.25125%
1.25125% of 4.0029g = 0.0500 (0.0125125 x 4.0029)
0.0500 + 4.0029 = 4.0529
4.0029 – 0.0500 = 3.9529
So therefore the confidence limit lies between 4.0529 + 3.9529
Forward titration (pure aspirin)
Balance error
Percentage error = 0.00005 x 100 = + 0.016%
0.3086
Pipette (class B) 25 cm3
Percentage error = 0.06 x 100 = + 0.24%
25
Burette (class B)
Percentage error = 0.05 x 100 = + 0.299%
16.725
Total percentage error = 0.016 + 0.24 + 0.299
= + 0.5632%
By working out the total percentage error, the confidence limit can be calculated.
Actual weight of aspirin = 0.3086g
Total % error = 0.5632%
0.5632% of 0.3086g = 0.0173 (0.05632 x 0.3086)
0.3086 + 0.0173 = 0.3259
0.3086 – 0.0173 = 0.2913
This means the mass of aspirin that reacted lies between the following values,
So therefore the confidence limit lies between 0.3259 + 0.2913
Forward titration (re-crystalised aspirin)
Balance error
Percentage error = 0.00005 x 100 = + 0.0165%
0.3032
Pipette (class B) 25 cm3
Percentage error = 0.06 x 100 = + 0.24%
25
Burette (class B)
Percentage error = 0.05 x 100 = + 0.3067%
16.3000
Total percentage error = 0.0165 + 0.24 + 0.3067
= + 0.5632%
By working out the total percentage error, the confidence limit can be calculated.
Actual weight of aspirin = 0.2934g
Total % error = 0.5632%
0.5632% of 0.2934g = 0.0165 (0.05632 x 0.2934)
0.0165 + 0.2934 = 0.3099
0.2934 – 0.0165 = 0.2769
This means the mass of aspirin that reacted lies between the following values,
So therefore the confidence limit lies between 0.3099 + 0.2769
Back titration (pure aspirin)
Balance error
Percentage error = 0.00005 x 100 = + 0.00332
1.5067
Pipette (class B) 25 cm3
Percentage error = 0.06 x 100 =+ 0.24%
25
Burette (class B)
Percentage error = 0.05 x 100 =+ 0.593%
8.4250
Volumetric flask (class B) 250 cm3
Percentage error = 0.2x 100 = + 0.08%
250
Total percentage error = 0.00322 + 0.593 + 0.24 + 0.08
= + 0.91622%
By working out the total percentage error, the confidence limit can be calculated.
Actual weight of aspirin = 1.4850g
Total % error = 0.9162%
0.9162% of 1.4850g = 0.1360 (0.05632 x 1.4850)
0.1360 + 1.4850 = 1.621
1.4850 – 0.1360 = 1.349
This means the mass of aspirin that reacted lies between the following values,
So therefore the confidence limit lies between 1.621 + 1.349
Back titration (re-crystalised aspirin)
Balance error
Percentage error = 0.00005 x 100 = + 0.003298%
1.5161
Pipette (class B) 25cm3
Percentage error = 0.06 x 100 = + 0.24%
25
Burette (class B)
Percentage error = 0.05 x 100 = + 0.588%
8.50
Volumetric flask (class B) 250 cm3
Percentage error = 0.2 x 100 = + 0.08%
250
Total percentage error = 0.003298 + 0.588 + 0.24 + 0.08
= + 0.911298%
Total percentage error = 0.00322 + 0.593 + 0.24 + 0.08
= + 0.91622%
By working out the total percentage error, the confidence limit can be calculated.
Actual weight of aspirin = 1.4850g
Total % error = 0.9162%
0.9162% of 1.4850g = 0.1360 (0.05632 x 1.4850)
0.1360 + 1.4850 = 1.621
1.4850 – 0.1360 = 1.349
This means the mass of aspirin that reacted lies between the following values,
So therefore the confidence limit lies between 1.621 + 1.349
Melting point apparatus
Percentage error = 0.10x 100 = + 0.10%
100
Total percentage error = + 0.10%
Preparation of 0.1 mol dm-3 sodium hydroxide solution
Balance error
Percentage error = 0.00005 x 100 =+ 0.0025%
2.0000
Volumetric flask (class B) 500 cm3
Percentage error = 0.25x 100 = + 0.005%
500
Total percentage error = 0.0025 + 0.005
= + 0.0075%
Preparation of 1.0 mol dm-3 sodium hydroxide solution
Balance error
Percentage error = 0.00005 x 100 = + 0.005%
10.0000
Volumetric flask (class B) 250 cm3
Percentage error = 0.2x 100 = + 0.08%
250
Total percentage error = 0.005 + 0.08
= + 0.0.085%
There are procedural errors that can occur if the practical technique is not good such as:-
- The solution made in the volumetric flask may not have been mixed thoroughly.
- The burette and pipettes may not have been washed with the solutions being used
- The conical flask may not have been washed thoroughly
- Over shooting of the end point may have taken place
- Too many drops of indicator may have been added.
- The meniscus may not have been on the graduation mark.
To avoid these errors the following techniques were implemented:-
- The solutions were added to the conical flask via pipettes. This delivered the specific volume accurately.
- The solution made in the volumetric flask was mixed thoroughly and the meniscus was accurately put on the graduation mark drop by drop by using a pipette. The solution was mixed thoroughly to make it homogenous.
- The burette was washed with distilled water first to clean it and then washed with the acid solution. The burette was washed with the acid to get rid of any water present in the burette and any air bubbles. If water was present in the burette it would make the acid weaker and therefore more acid is needed to neutralise the sodium hydroxide. If air bubbles were present it would affect the final reading of the burette.
- The pipette was firstly washed with distilled water to clean it. It was then washed with the sodium hydroxide to get rid of any water that may be present in the pipette. If water was present in the pipette it would have made the sodium hydroxide solution more dilute and therefore more acid would have been needed to neutralise it.
- The conical flask was washed thoroughly with distilled water before each titration to get rid of any previous solution present.
- A rough titration was carried out to give an approximate value of acid added to the flask to reach the endpoint. The titrations were repeated to obtain two or three concordant results and then an average was calculated.
- To avoid overshooting of the endpoint, acid solution was added from the burette drop by drop. The conical flask was swirled to mix the solutions thoroughly when the acid was being added. Care was taken when doing so, so that no liquid was lost through splashing.
- Three drops of indicator were added for each titration to make it a fair investigation, because different amounts of indicator give different endpoints.
Other experiments could be carried out such as colorimetry, thin layer chromatography and Iron (III) chloride test.
Acknowledgements
Many thanks to my project supervisors Dr Knutton and Mr Gallagher at the Huddersfield Technical College for supporting me throughout my project.
I would also like to thank Sheila our Technician at the Huddersfield Technical College for preparing all the apparatus, chemical and solutions for my project.
References
(1)
(2) Chemical storylines 2nd edition pg113-116
George Burton, John Holman, John Lazonby, Gwen Piling, David
Waddington
(3)
(4)
(5)
(6)
(7) What’s in a medicine 3rd edition WM5.1
(8)
(11)http://courses.cm.utexas.edu/pmccord/spring2005/ch455/Spr05455Wk2Lab.pdf
(12)http://chemlabs.uoregon.edu/Classes/Exton/CH229/AspAnalysis/AspAnalysis.pdf
(13) What’s in a medicine 3rd edition WM6
(14)
(15)
(16)http://en.wikipedia.org/wiki/Back titration
(17) Chemical ideas 2nd edition pg12
George Burton, John Holman, John Lazonby, Gwen Piling, David
Waddington