Candidate code: jyj422 (001307 -0014)
Topic: Kinetic Study of Synthesis
Research Question: What is the effect on the rate of using different acids as a catalyst for esterification
reaction?
Introduction
Esterification, as the name suggests, is the process of creating esters. These organic compounds are
widely used in flame retardants, gasoline, oil additives, and perfumes (Helmenstine, 2019). The use of
esters in perfumes is what separates it from most other chemical compounds and proves the importance of
these compounds in our daily lives. As esters are an essential component of several products, many
industries require enormous amounts of these compounds. However, as esterification is a slow reversible
exothermic reaction (Azhar et al., (n.d)), industries compromise their equilibrium conditions to produce
the maximum amount of ester in the least possible time. Although an exothermic reaction, heat is provided
coupled with a concentrated H2SO4 catalyst to increase the reaction rate for esterification. (Azhar et al.,
(n.d))
The catalyst simultaneously lowers the activation energy (Ea). Increasing the temperature increases
the average kinetic energy of the reactant molecules, so more molecules acquire Ea, resulting in greater
successful collisions. Catalysts can accelerate a reaction without polluting its results. It was surprising that
industries only used concentrated sulfuric acid as a catalyst for esterification, but it was unclear why.
However, as part of school policy, all SL students must attend HL classes, and in one class, the teacher
discussed electrophiles and nucleophiles. It was discussed that esterification is a nucleophilic substitution
reaction that only requires an H+ ion, which any acid can provide. Therefore, why is only sulfuric acid used
in industrial esterification and not any other concentrated acid? To answer this question, this experiment
evaluates the effects and reliability of 3 concentrated acids as catalysts in the esterification reaction for 2propyl ethanoate to determine whether there is a better catalyst than the conventionally used one.
Background:
As esters are organic compounds, they experience specific arrangement patterns that divide them
into functional groups. Functional groups are important as they are critical in defining the reaction
(Libretexts, 2020). The ester functional group forms due to substituting the alcohol and carboxylic
functional group. The alcohol functional group involves a single bond between a carbon and an OH (also
called hydroxyl) (Libretexts, 2020). The hydroxyl group reacts with the carboxyl group (COOH) from the
carboxylic acid to form the ester functional group and water. As esterification is a slow process, it requires
an acid catalyst to lower the activation energy of the reaction and assist in the nucleophilic substitution
reaction between the hydroxyl and carboxyl groups, as shown in Fig. (1).
The carboxyl group is protonated by an H+ ion from the H2SO4 catalyst as oxygen shares an electron
with it; oxygen gains a positive charge. As oxygen is more electronegative than carbon, it steals one of the
double C=O bond to eliminate its positive charge. However, a positive charge is formed on C, leaving it
susceptible to a nucleophilic attack from the hydroxyl group. The hydroxyl oxygen shares an electron with
the positive C and acquires a positive charge. As oxygen is more electronegative than H, it attracts the OH to balance the positive charge, leaving an H+ ion. This H+ ion receives an electron from the oxygen in
the OH in carboxyl, forming a positive charge on that oxygen. The positive oxygen atom of the carboxyl
pulls the C-O bond to balance its charge and leave a positive charge on C. The oxygen, therefore, forms
H2O, and the elimination of water occurs. The positive C atom attracts a bond from the oxygen, forming a
C=O bond and leaving a positive charge on O. The positive oxygen pulls an electron from its O-H bond,
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and deprotonation occurs. The released H+ bonds with the OSO3H- to form H2SO4. Hence, esterification is
complete, and the catalyst does not pollute the products.
Fig. (1): Nucleophilic substitution reaction of esterification.
(Khan Academy. (n.d.))
The ester functional group consists of a carbon bonded to 2 oxygen. As shown above, there is a
double bond between the carbon atom and one oxygen atom and a single bond between the carbon atom
and the other oxygen atom. The parent alcohol and carboxylic acid react together by extracting the water
molecule and combining the two reactants (Libretexts, 2021). Once the ester is formed and extracted, it has
several applications such as perfumes, preservatives, toothpaste, hair-care products, cosmetics, deodorants,
moisturizers, etc. Varying the reactants in esterification can help obtain these unique properties (BYJU'S,
2021). Still, the industries producing these esters need to increase the rate of reaction as much as they can
and to do so, they use catalysts. However, not all catalysts affect the rate of reaction in the same way, and
it is crucial to use the right one to minimize the reaction time.
Hypothesis:
Ha: The reaction will occur but will be slower with hydrochloric and nitric acids compared to
sulfuric acid.
H0: The reaction will occur but can be faster for nitric and/or hydrochloric acid and will be faster
compared to sulfuric acid.
Apparatus:
1.
2.
3.
4.
5.
6.
7.
Stirring rod
Hot water bath.
Bunsen Burner
PPE (safety glasses, safety shield, nitrile gloves, neoprene aprons)
Thermometer
dropper
Pipette
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8. Test tubes
9. Beaker
10. Test tube rack
11. Tripod claw for a thermometer
12. Tripod stand for the water bath
Reactants Required:
Purely chemical Ethanoic acid (CH3COOH)
Purely chemical 2-propanol (C3H8O)
Sulfuric Acid (concentrated)
Hydrochloric acid (concentrated)
Nitric Acid (concentrated)
Variables:
Independent variable
Dependent variable
Controlled variables
Type of catalyst used.
Rate of change in temperature
The concentration and volume of
the reactants
2-propanol and ethanoic acid
(reactants).
Amount of catalyst (5 drops)
Procedure:
1) Prepare a water bath at 70°𝐶and maintain that temperature (70°𝑪 is an elevated temperature and
could cause burn; therefore, manage hot apparatus with caution).
3) Using a pipette, take 5.0 ml of alcohol in a test tube.
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4)Using a pipette, take 10.0 ml of ethanoic acid in a separate test tube.
5)Using a dropper, take five drops of catalyst and add it to the alcohol in the test tube (Use PPEs when
dealing with concentrated acids).
6)Place both test tubes in the water bath and monitor the temperature of each liquid using separate
thermometers (As water would be near boiling point, be careful when transferring the test tubes
as any contact could cause burns).
7) Wait until the temperature of all 3 test tubes is equal to the water bath at 70°𝐶.
8) Pour the alcohol and catalyst test tube into the acid test tube within the bath (use clamps to hold
the test tubes as they would be hot).
9)Monitor temperature using a thermometer.
10)Record the time taken to reach the maximum temperature.
11)Find the temperature difference between the maximum temperature and the controlled temperature
of the bath.
12)Repeat the experiment with different catalysts.
Methodology:
The reaction rate is usually calculated using the mass of the reactant consumed per unit time and
the mass of product created per unit time. If this is not possible, color change, electrolyte difference, or
volume of gas produced are used. Since both the reactants in this experiment are solvents and produce no
precipitate, the reaction rate must be measured through the invisible product: the heat of reaction. As
esterification is an exothermic reaction that produces heat as a product, the rate of temperature change
would help analyze how much product is produced per unit time, therefore measuring the reaction rate due
to varying catalysts.
However, the reaction is slow, so heat is required to increase the rate of reaction in addition to a
catalyst. To provide heat, usually, a rheostat would be used; however, as the lab’s equipment was not so
updated, a hot water bath was used to heat all the reactants to the desired equal temperature, which would
provide the particles enough kinetic energy to overcome the activation energy and react successfully. A
half-filled metal pot with water was placed on a lit Bunsen burner which heats the water to the desired
temperature of 70°𝐶. Whenever during the experiment, the temperature rose above 70°𝐶, the heat was
turned off to maintain the temperature, allowing the water bath and the reactants to remain at a constant
temperature so that when the reactants are later mixed, the temperature change recorded by the thermometer
was only due to the heat released because of the reaction and not because of the difference in temperature
of the reactant. After maintaining the temperature of the water bath, the reactants were measured according
to steps 3, 4, and 5 of the procedure. The test tubes were placed in a test tube holder rack and placed in the
water bath.
Once the temperatures of the reactants were constant at 70°𝐶 step 9 was conducted while keeping
the ethanoic acid tube submerged in the water bath and the thermometer still in place. Otherwise, the
temperature reading would have dropped if the test tubes were pulled out and mixed. This would have led
to inaccurate temperature readings as there would have been a temperature difference between the reaction
system and the hot water bath. So, the increase in temperature due to the transfer of energy would have
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overshadowed the increase in temperature due to the exothermic reaction. Mixing the reactants while
keeping the ethanoic test tube in the water bath kept the reactants at a constant temperature. The focus was
solely on the increase in temperature due to esterification.
Once the reactants were mixed, the temperature change was measured (temperature change above
the temperature when the reactants were mixed). The maximum temperature reached and the time taken
were noted for calculations. The reaction had occurred, and there was 2-propyl ethanoate, water, and excess
ethanoic acid present in the solution. The experiment was finished at this stage because now the ester would
have been extracted in esterification. However, the final crude ester was not separated from the excess acid
as the experiment aimed to record the reaction rate through the change in temperature due to different
catalysts to ensure the most effective one.
Raw data:
Catalyst: H2SO4
Catalyst: HCl
Time
(hours: minutes)
(±0.1 ms)
Temperature
±0.5°C
0:00
(test tubes added
to the water
bath)
75.0°C
0:05
70.0°C
0:10
69.0°C
0:15
70.0°C
(reactants
mixed now)
instant rise to
71°C
0:20
70.0°C
0:25
70.0°C
Time
(minutes: seconds)
(±0.1 ms)
Temperature
±0.5°C
0:00
(test tubes added to the
water bath)
70.0°C
0:05
72.0°C
0:10
70.0°C
(reactants mixed
and instant rise to
71.5°C)
0:15
71.5°C
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Candidate code: jyj422 (001307 -0014)
Catalyst: HNO3
Time
(minutes: seconds)
(±0.1 ms)
Temperature
±0.5°C
0:00
(test tubes added to the water bath)
72.0°C
0:05
70.0°C
0:10
70.0°C
(reactants mixed
and instant fall to
69.0°C)
0:15
70.0°C
Processed data:
Balanced Esterification Equation:
C3H8O + CH3COOH → CH3COOCH2CH3 + H2O
Step#1: Calculating average bond enthalpies:
Average Bond Enthalpies:
Reactants:
Propan-2-ol:
Ethanoic Acid:
2(C-C) = 692
3(C-H) = 1242
7(C-H) = 2898
(C=O) = 804
(C-O) = 358
(C-O) = 358
(O-H) = 463
(O-H) = 463
Average Bond enthalpy of reactants = 7624 kJ/mol
Products:
2-propyl ethanoate:
H2O:
10 (C-H) = 4140
2(0-H) = 926
3(C-C) = 1038
(C=O) = 804
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2(C-O) = 716
Average Bond enthalpy of products: 7624 KJ/Mol
Enthalpy change: 7624-7624 = 0 (according to average bond enthalpies given in the data booklet)
This goes against the theory that an enthalpy change of reaction value would be calculated through the
bond enthalpies. However, it was learned that the bond enthalpy values are an average of the bond across
all different chemical substances. Therefore, if the enthalpy change of reaction value is minimal, this
inaccuracy of the bond enthalpies may provide such answers. Hence, research was conducted to find the
enthalpy change of this reaction, and this value was obtained:
Heat of this reaction (𝛥𝑟 𝐻 0 ) = −1.289 kJ/mol (Mekala, 2021)
Hr
Step#2: Calculate the moles of acid and alcohol used.
Moles of Alcohol used:
𝑑 = 𝑚/𝑣
d= 786 kg/m3 =0.786 g/ml (density of alcohol) (Fleming, 2020)
The Volume of alcohol= 5ml
Mass = (Density * Volume) = 0.786 * 5 __________________ eq. (i)
= 3.93 g
Sol. (i)
Using Sol. (i): Actual Mass = Percent Purity * Total Mass = 3.93g * 0.99______________ eq.(ii)
= 3.89g
Sol. (ii)
Mr (2-propanol) = 60.1g/mol
Using Sol.(ii): moles of 2-propanol = 3.89/60
eq. (iii)
= 0.0648 mol. of alcohol _________________________________Sol. (iii)
Moles of Acid used:
d= 1.05g/cm3= 0.00105 kg/ml (density of ethanoic acid) (Acetic acid (n.d))
Volume of ethanoic acid= 10ml
Mass= (Density*Volume) = 1.05 * 10__________________________eq. (iv)
10.5 g
Sol. (iv)
Using Sol. (iv): Actual Mass = Percent Purity * Total Mass = 10.5 * 0.99 ____eq. (v)
= 10.395g _______________________________________________________Sol. (v)
Mr (ethanoic acid) = 60.052 g/mol
Using Sol. (v): moles of ethanoic acid= 10.395g/ 60.052 g/mol______________ eq. (vi)
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= 0.173 mol. of ethanoic acid_________________________________________ Sol. (vi)
C3H8O + CH3COOH → CH3COOCH2CH3 + H2O
Mole ration of Alcohol to Acid is 1:1
Comparing Sol. (vi) and Sol. (iii), it can be determined that 0.0648 moles of acid are required (as 0.173
moles of acid were available, acid was in excess, and the alcohol was the limiting reagent which determines
how much ester is produced.)
Using Sol. (iii): Alcohol: Ester = 1: 1 ________________________________ eq. (vii)
0.0648 mol. of ester _____________________________________________Sol. (vii)
Step#3: Calculate the enthalpy change of reaction:
Using Hr from step#1 and Sol. (iii) from step#2:
Mass of reaction mixture: 15g
Theoretical energy released:
𝛥𝑟 𝐻 0 = Hr / no. moles
-1.289 kJ/mol*0.0648
Heat= 0.08359 kJ (theoretical value)
eq. (viii)
Sol. (viii)
Step#4: Theoretical temperature change for esterification reaction of 2-propyl ethanoate
𝑄 = 𝑚𝑐𝛥𝑇
𝑄 = 0.08359 𝑘𝐽
(Energy value taken from Sol. (viii)) (Engineering ToolBox. (n.d.))
Mass= 0.014285 kg (sum of Sol. (ii) and Sol. (v)) (Theory of conservation of mass)
Specific Heat Capacity of Ethanoic acid= 2.043𝑘𝐽/(𝑘𝑔𝐾) (ScienceDirect, (n.d.))
Δ Temperature=?

As acid was in excess, it is assumed that there would be ester and excess ethanoic acid present in
the mixture after the reaction. As excess ethanoic acid would have more moles than the ester, most
of the mixture would be ethanoic acid. Therefore, it can be assumed that the specific heat capacity
of the mixture is the same as ethanoic acid.
0.08359
ΔT= (0.014285)(2.043) = 2.7°𝐶

Sol#8
As the reactants were the constant variable, the value in Eq#7 can be assumed as the theoretical
temperature change value for the three trials conducted in this experiment.
Step#5: Comparing the experimental results with the theoretical value of Eq#7:
Theoretical temperature change = 2.7°𝐶
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Temperature increases with catalysts:
Concentrated H2SO4 = 1°𝑪
Concentrated HCl = 1.5°𝑪
HNO3 = -1°𝑪
Safety Precautions:
As concentrated acids were used in this experiment as catalysts, safety precautions were crucial. It was
essential to wear a lab coat, medical face mask, protective gloves, eye protection, and footwear that
completely covered the feet throughout the experiment. Moreover, this experiment had to be conducted
under the supervision of the chemistry teacher and the lab in charge due to the hazardous chemicals used.
The concentrated acids had to be managed only by the lab in charge as per safety precautions. The lab in
charge looked after the school lab for several years and was qualified as per the IB guidelines to manage
corrosive chemicals.
Moreover, the temperature of the water bath had to be maintained, which required constantly turning the
Bunsen burner on and off. Therefore, the lab in-charge’s assistance was utilized when dealing with hot or
flammable equipment. The test tubes within the water bath were also managed by the lab in charge when
mixing the reactants. The chemical bottles had to be capped as soon as used to prevent fumes from filling
up. The apparatus that was in contact with the reactants were thoroughly rinsed, and the equipment used
for the catalysts was rinsed by the lab in charge, as the lab in-charge entirely managed the concentrated
catalysts.
Error / Uncertainty:
Thermometer = ±0.5°𝐶
(Variable uncertainty as systematic uncertainty was proved to be zero by
the steam point and ice point method).
Pipette = ±0.02 𝑚𝑙 (Stone, 2008)
Stopwatch = ±0.1 𝑚𝑠
Evaluation:
A hot water bath was set up to provide heat to the reaction as expressed in the methodology; however,
initially, it was decided to set the temperature at a constant of 80°𝐶. While experimenting, the boiling points
of 2-propanol and ethanoic acid (reactants) were checked. It revealed that the boiling point of 2-propanol
is 82.5°𝐶. Therefore, it was decided to lower the constant temperature of the water bath by 10°𝐶. As the
temperature of the water bath was controlled by manually turning the burner on and off, the temperature
of the water could potentially have reached the boiling point of the alcohol, which would have lowered its
volume in the reaction. At 70°𝐶, all the reactants would have enough energy to react faster, especially
coupled with the effect of the catalyst.
Moreover, once the water bath had reached the constant temperature, the test tubes were added, and
thermometers were placed in both the test tubes to monitor their temperature. Surprisingly, the alcohol and
catalyst test tube were heating slower than the ethanoic acid. Initially, it was assumed that this difference
was due to differences in specific heat capacities. After a while, the ethanoic acid was at the same
temperature as the water bath. However, the alcohol and catalyst were still much below 70°𝐶. After waiting
for a few more minutes, it was suspected that there could be a problem with the thermometer, so another
thermometer was added to the alcohol test tube. Soon, it was evident that the previously used thermometer
was inaccurate as the new thermometer was giving a reading much closer to the 70°𝐶 mark. In contrast,
the previous one was still displaying a lower temperature. Therefore, it was decided to calibrate the lab’s
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thermometers using the steam point and ice point method (Raiseupwa.com (n.d)). After assessing four
thermometers, it was discovered that three were accurate and one was inaccurate. The three were giving
completely accurate readings. For the ice point experiment, they precisely showed 0°𝐶. For steam point,
they precisely showed 100°𝐶. Therefore, the systematic uncertainty was zero when measuring the
dependent variable; however, there was still a possibility of random errors. After having the lab in charge
inspect the old thermometer, he confirmed that the thermometer bulb was broken.
The experiment was restarted with the correct and accurate thermometers. After waiting for a bit longer,
the reactants were mixed when the entire system was at 70°C. It was decided that the test tubes would be
lifted from the water bath, and the contents of one would be poured into the other. The lab in-charge’s
assistance was utilized to take out the test tubes and pour the alcohol and catalyst into the acid test tube.
However, as mentioned in the methodology, it caused the temperature to fall below the temperature of the
water bath, which could affect the temperature reading and temper the results. As the required quantity was
only the temperature change due to the esterification reaction, this method would not have been effective.
The temperature rise would have also been due to the difference in the temperature of the water bath and
the reaction. Therefore, the procedure was modified. The alcohol and catalyst solution were poured into
the ethanoic acid solution while keeping the ethanoic acid test tube in its place in the water bath. This would
keep the temperature of that test tube at a constant level equal to the water bath, and any increase in
temperature observed would be due to the exothermic reaction.
The experiment was repeated with the updated methodology, and the results were surprising. The
temperature increase observed was around 1 ⁰C and occurred instantaneously. This was because of the
lesser amounts of reactants used. Also, the temperature of the alcohol was near its boiling point, so the
particles had enough kinetic energy. This high kinetic energy coupled with the catalyst allowed the reaction
to occur instantly. As the reaction was exothermic, an increase in the thermometer’s temperature reading
would indicate the product being formed. The greatest temperature change was expected for the
concentrated sulfuric acid as it is considered the best catalyst in the industry for esterification; its reaction
should have produced a higher reading than the other catalyst. However, after conducting the experiments,
the HCl reaction recorded a maximum temperature of 71.5⁰C, whereas the H2SO4 reaction recorded only
71⁰C. Therefore, according to the experiment, HCl is the more efficient catalyst as it produces more heat,
one of the products of an exothermic reaction. The results were confusing. Therefore, I further researched
catalysts in esterification and why HCl is not used in place of sulfuric acid.
The research revealed that esterification reactions produce ester, heat, and water. The heat is emitted, and
the ester and water remain in the solution, of which only the ester is required. Concentrated sulfuric acid is
used because it acts as a dehydrating agent and a catalyst. As a result, it helps remove the water from the
solution, which concentrated HCl or concentrated HNO3 could not do. Moreover, the Cl- ions can take part
in any side reactions with the alcohol, polluting the results.
Limitations:
As the reactants had to be at the same temperature, they had to be heated. However, the water bath created
relied on a Bunsen burner, and the temperature had to be monitored using a thermometer. Manually trying
to control the temperature could have affected the temperature readings and the reaction rate. An electric
heater or a rheostat that could maintain the temperature at 70⁰C could eliminate this limitation. The
temperature was recorded using analog thermometers which could be subject to human error. One of the
thermometers used initially was broken, which caused the first trial to fail. The thermometer was changed
and made sure it was working perfectly. Moreover, the system was not insulated, allowing most of the heat
to escape, resulting in lower thermometer readings. Using a digital thermometer would produce more
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accurate results. Overall, the apparatus used was unsuitable for esterification reactions, and other available
apparatus had to be used to make this experiment work.
Moreover, this experiment was based on a few assumptions, such as in step#4 that the specific heat capacity
of the ester is the same as the acid as there was excess acid left after the reaction and that the temperature
change recorded by the thermometer would only be due to the formation of ester. These assumptions might
have increased the absolute error as the volumes were minute. Moreover, small amounts of reactants were
used, and each catalyst was only tried once. The small volumes can lead to a higher overall percentage
error even for a very small uncertainty. This limitation occurred as small volumes are preferred in reactions
where heat changes are being measured as it does not dilute within large volumes.
Conclusion:
The alternate hypothesis has been confirmed as a catalyst does affect the rate of reaction, and each catalyst
produces a different rate of reaction. Surprisingly, HCl caused a higher reaction rate than H2SO4. These
results were confusing, and triggered further research into esterification to explain why H2SO4 is the
conventional catalyst. H2SO4 is di-protic, which allows a higher concentration of H+ ions for nucleophilic
substitution. H2SO4 is also a dehydrating agent and has the Le Chatelier's effect as it removes the water
from the reaction. Therefore, as the reaction is reversible, removing some of the product shifts the
equilibrium to the product’s side, allowing more of the ester to be produced. Moreover, with the water
gone, the ester can be separated from the mixture by neutralizing the excess acid and using fractional
distillation to separate the final ester. In conclusion, the effect from each catalyst was nearly the same
however, H2SO4 is the best catalyst for esterification due to being di-protic and a dehydrating agent, as
water is one of the products of esterification and removing water from the solution provide the ester.
Appendix:
Ice Point:
Steam Point:
Thermometers:
Safety Precautions:
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