Title: Kinetics and Aspirin Hydrolysis Composite Lab Report
This experiment included two ester hydrolysis reactions. A hydrolysis reaction refers to the dissociation of water. The first hydrolysis focused on the kinetic aspect of esters in acidic conditions. Methyl ethanoate was hydrolysed under acidic conditions during the experiment, thus enabling its concentration to be ascertained at regular intervals through titrations. The linear trend in both tabulated and graphical data ascertained follow first-order kinetics. The latter part of the experiment involved basic hydrolysis of aspirin. The sample was analysed to ascertain its purity, melting point and functional groups.
The aim of this experiment was to determine the rate of a reaction. The rate of reaction can be ascertained by calculating the concentration changes of the products at 10-minutes intervals over a total period of 90 minutes, highlighting the kinetic properties of the reaction.
The speed at which a reaction takes place can be determined by the rate of the reaction; making it an important aspect of pharmacokinetics. In addition, certain drug properties; shelf life and elimination rate can also be determined through rate of reaction.
Furthermore, another aim of the experiment was to ascertain a rate constant (k). The rate of reaction was directly proportional to the concentration and thus, we can assume the reaction followed first-order kinetics. Ethanoic acid and methanol are formed. Appendix 1 shows the mechanism of the experiment. The equation is:
The results of the experiment produced a straight line graph from plotting the ln(T? – Tt) of the titre values obtained at each interval, versus time (t). Therefore, the experiment followed first-order kinetics
The latter experiment involved the hydrolysis of aspirin using NaOH. Saponification is known as the basic hydrolysis of esters. NaOH was used to maintain basic conditions and to simulate alkaline conditions, or the location at which the reaction would take place; in the gastrointestinal tract. Appendix 2 shows the mechanism for the reaction and the equation is:
Aspirin is a non-opioid analgesic, also known as acetylsalicylic acid. The medication is commonly used to treat moderate pains, aches, fever and inflammation.
The pharmacodynamics and kinetics of the drug were investigated during this experiment. This important as esterase enzymes hydrolyse aspirin in the body into salicylic acid. The hydrolysis of aspirin can also occur if the drug is stored in humid conditions. Analysing the drug involved using infrared (IR) spectroscopy to determine functional groups, percentage yield to determine purity, and melting point; analysed against literature values.
Safety and COSHH:
Sodium hydroxide – Irritant, flammable, irritant
Methyl ethanoate – Flammable, irritant, harmful
Hydrochloric acid – Corrosive, irritant
Phenolphthalein – Flammable, irritant
Special Requirements / Other: Wipe up spillages, wash solution off skin with a copious amount of water. Wear lab coat and safety glasses always. Wear gloves when handling the HCl.
Disposal: Dispose Methyl Ethanoate in fume cupboard.
Aspirin – Harmful
Sodium Hydroxide – irritant, Corrosive
Sulfuric acid – Corrosive, toxic, irritant
Gloves worn. Followed lab safety precautions.
Special Requirements / Other: Wear gloves when measuring Sodium Hydroxide solution and concentrated Sulphuric Acid. Wipe up spillages immediately, using 20% sodium carbonate for any spillages of concentrated acid. Wash splashes off skin with a copious amount of water.
Wear lab coat and safety glasses at all times.
Disposal: Filter paper and products to be disposed of in fume cupboard.
A 250ml conical flask was labelled ‘Tt’. 50ml of distilled water was measured using a measuring cylinder and placed into the flask. 50ml of 2M HCl was then pipetted into the conical flack using a 50ml volumetric pipette. A stopper was placed in the flask and was then immersed in a water bath, at 25C. A timer was set to 10 minute, to allow the solution to equilibrate.
250ml conical flask was labelled ‘T?’. 50ml of distilled water was added to the flask, again, using a measuring cylinder. 5ml of methyl ethanoate was then also added to the flask. The flask was then immersed in a water bath for 60 minutes to allow it to reach equilibrium.
Once flask ‘Tt’ had time to equilibrate. A 50ml burette was filled with
0.2 M NaOH. A 5ml volumetric pipette was used to pipette 5ml of methyl ethanoate into flask ‘Tt’. The flask was then shaken thoroughly and a stopwatch was started for 10 minutes. The flask was kept in the water bath, still at 25°C. After the stopwatch had reached 10 minutes, a 5ml aliquot of flask ‘Tt’ was pipetted into a 250ml conical flask containing 50ml of ice cold distilled water.
Immediately, the sample was then titrated against the 0.2 M NaOH with 3 drops of phenolphthalein indicator. The colour change from colourless to a light pink solution indicated the end-point of the titration. The titre value was recorded in a table. This process was then repeated every 10 minutes over a 90-minute period.
Once the 90 minute period was complete, 50ml of distilled water was measured using a measuring cylinder and placed into a 250ml conical flask. A 5ml aliquot of flask ‘T?’ was taken and pipetted into this flask using a 5ml volumetric pipette This was then titrated against 0.2 M NaOH, again using 3 drops of phenolphthalein indicator. Again, the colour change from colourless to a light ink solution indicated the end-point was reached. Concordant results were obtained by repeating this titration. The results of this were again, recorded in a table.
A 0.2641g sample of aspirin was accurately weighed by difference using a weigh boat on an analytical balance. The sample was then transferred to a clean test tube and 10ml of NaOH was added to the sample, using a measuring cylinder. The test tube was mixed by swirling the test tube until the entirety of the solid had completely dissolved. The test tube was then placed in a steam bath, using tongs to boil the solution for approximately 5 minutes, in order for the reaction to complete. The solution was then cooled to room temperature. Once the solution had cooled, concentrated sulphuric was carefully added a few drops at a time, until the solution turned acidic. Acidity was identified by using a stirring rod to transfer a small drop of the solution onto the litmus paper. The litmus paper turned red when acidity was reached. Using Büchner filtration apparatus, the resulting precipice was filtered. The content of the test tube was washed with distilled water to ensure as much of the precipitate was transferred onto the filter paper as possible. After roughly 10 minutes, the distilled water was removed from the precipitate. A vial was then weighed using an analytical balance. Using a spatula, the filtrate was then transferred into the vial. The percentage yield was then calculated. The melting point of the filtrate was then obtained using melting point apparatus and the melting point obtained was compared to the known literature value of 158-161°C, in order to evaluate purity. The filtrate was then analysed using IR spectroscopy, allowing the identification of the functional groups present.
Titre at time t, Tt
T? – Tt
ln (T? – Tt)
Figure 1 – Figure 1: Table to show the titration results for flask T? against 0.2 M NaOH.
Flask titre (T?) = 40.1 mL
Figure 2 – Graph of ln (T? – Tt) versus time.
Product (salicylic acid)
Weight of vial (g)
Weight of sample + vile (g)
Weight of sample (g)
Figure 3 – Table to show weigh by difference results of Aspirin and Salicylic Acid.
Percentage yield calculation:
Moles of Aspirin = Mass / Mr
= 0.2641 / 180.16
= 1.47 x 10-3 mol
Mass Salicylic Acid = Moles x Mr
= (1.47 x 10-3) x 138.12
Percentage yield = (Actual yield x 100) / Theoretical yield
= (0.1071 / 0.202) x 100
Sample Initial Melting Point: 159.5oC
Sample Final Melting Point: 161.2C
Literature Melting Point: 158oC – 161oC (British Pharmacopeia)
Figure 4 – IR spec of salicylic acid sample produced.
To verify the accuracy of my rate constant value, 1.24×10-2min-1, it was substituted into the equation for a straight line, y=mx+c. The y value can be ascertained as the intercept of the graph and the x value can be a time value. A first-order reaction is also demonstrated in the graph as there is a clear negatively correlated, linear line.
Heating the two flasks was vital in ensuring that the full hydrolysis reaction had taken place. To obtain an accurate sample of ethanoic acid and to prevent further hydrolysis from happening, ice cold water was used to slow down the reaction. The method of quenching was applied through using the ice-cold water. Appendix 1 demonstrated the mechanism of reaction. As the graph did not produce an exact linear trend line, we can assume that despite using the quenching method, it still allowed room for anomalies. Thus, a line of best fit was drawn to ascertain the rate constant, k.
A possible reason for such anomalies could be a result of the temperature of the water. If the temperature was not low enough, the reaction would not have stopped and the reaction would have continued to take place during the titration. As a result of this, an inaccurate volume for the tire of ethanoic acid would have been produced. Another possible cause for the presented anomalies could be due to human error in judging exactly when the colour change; the end-point took place. It is difficult to equally judge this for each interval using the human eye only.
As the stomach contains HCl, many drugs containing the common functional group ester, are susceptible to undertake acid hydrolysis reactions. The dissolution rate and thus acid hydrolysis of some drugs in the stomach, including aspirin, may cause discomfort and irritation to a patient. Due to this, some tablets are given an enteric coating that is acid resistant and thus prevents dissolution in the stomach; preventing discomfort (The International Agency for
Research on Cancer, 1997). Dissolution is however, in non-acidic conditions achieved using this coating, normally in the small intestine where the drug can be absorbed.
The determination of the rate of the experiment was the main of the experiment, in which an acid catalysed an ester hydrolysis. Demonstrated in my results in figure 2, it is clear that the experiment was successful as the titre values increased with an increase in time. This is most likely due to the volume of NaOH required to obtain the end-point was directly proportional to the volume of ethanoic acid produced by the hydrolysis.
The principal aim of this experiment was to use base hydrolysis of aspirin to analyse a sample of salicylic acid to ascertain its purity. The calculated percentage yield was only 53.02%. This is extremely low and indicated the incompletion of the reaction. Multiple factors could have affected the percentage yield including the reaction of NaOH and aspirin not fully reacting or an insufficient amount of sulfuric acid added. The poor yield highlights the importance of allowing initial reactions to completely take place and to add acid in drops, in excess to help achieve the end-point.
Through analysing the infrared spectrum in figure 4, similarities were shown compared to the infrared spectrum of salicylic acid.
Although the fingerprint regions that are below 1500cm-1 of both spectra are similar (Maquelain, 2002), they provide little knowledge of the compound. However,
The infrared spectrum of the sample shown in figure 4, demonstrated characteristic vibrational peaks at wavenumber 323.77 cm-1 and 1986.68 cm-1 that were assigned to OH and C-H stretching, respectively. The C=O (COO-) asymmetric and symmetric stretching were assigned to infrared peaks observed at 1608.63 and 1382.96 cm-1, respectively. Further, infrared peaks appeared at 1577cm-1 were attributed to C=C multiple peaks. The C-C stretching peak was observed at 1440.83-. This suggests the sample was fairly pure, (Trivedi MK, 2015) despite the low percentage yield. The peaks appear to be slightly hindered in that they are not as intense as the sample spectra. This is most likely a result of small impurities in the sample and/or that the reaction had not fully complete.
Moreover, there also appears to be peaks on the infrared spectrum that are not present on the sample spectrum. This is again, most likely due to the presence of impurities.
The melting point of the sample was accurate (Joint Formulary
Committee, 2017)and was within the range stated in the literature values. This further suggests that the sample was fairly pure, as the presence of impurities would be likely to hinder and intermolecular forces by weakening them, making the melting point lower.
The kinetics of reactions and the calculations involved in rates of reaction were demonstrated through the experiment and an understanding of acidic and basic reactions was given. The reaction in the first experiment was first-order and the rate constant was 0.0124min-1. This was confirmed through substituting the y-intercept and rate constant into the equation for a linear graph. A straight-line trend which accentuated first-order kinetics was plotted on the graph through ascertaining a natural log concentration, by entering a time value. By calculating percentage yield, melting point and using infrared spectroscopy to identify functional groups, the purity of the salicylic acid; the product of the aspiring hydrolysis, the purity was given.
Although multiple, similar functional groups could be identified by the infrared spectrum, the infrared spectrum did not match completely, implying that the purity may not be as high as anticipated.
Appendix 1: Mechanism for hydrolysis of methyl ethanoate using HCl
Appendix 2: Mechanism for hydrolysis of aspirin using NaOH