Dawson Community College Chemistry Lab Report

HEATS OF REACTIONCHMY 141-220
TA: Olivia Andrus
11/14/19
Introduction
As the title indicates, this lab explored the properties of the heat released from certain
reactions through recording the heat change as the reaction occurs with the use of a thermistor.
Performing these reactions demonstrated some principles of thermodynamics; including the use
of Hess’s Law, which states that the total heat change of a reaction is equal to the sum of the heat
from all the changes in the reaction (Cohen, 2019). In this lab, three different reactions were
performed: two of the reaction’s heats should, theoretically, add up to be the heat of the third
reaction, precisely demonstrating the principles of Hess’s Law. The heat of each reaction was
recorded by setting up a thermistor suspended over a styrofoam cup that was filled with the
substance that was being reacted; the heat registered by the thermistor was then recorded
graphically on a computer program and used for interpretation. Thermodynamics and Hess’s
Law can be applied to many things, from car companies, to caloric breakdown of food in the
body. Most importantly, Hess’s Law is often used for industries that use the burning of fuel for
energy; with Hess’s law, the companies can measure how much energy a fuel releases and use
the information to make efficient energy choices (Harshini, et. al. 2015).
Procedure and Observations
This lab began by setting up a Microlab experiment with a thermistor that would record
the time and temperature of the chemicals being tested. A graph was prepared using time as the
x-axis, and temperature in Celsius as the y-axis. The thermistor was set up so it was suspended
over a styrofoam cup about one-half inch from the bottom. Next, the mass of a 50 mL beaker
was taken and recorded. Two to three grams of KOH pellets were measured out and the exact
mass was taken and recorded, then placed into the beaker, and the mass of the beaker and the
pellets were taken and recorded. This beaker was set aside for the moment and 100 mL of water
was measured using a graduated cylinder and poured into the styrofoam cup. The temperature of
the water was recorded for two minutes, then, while continuing to record, the KOH pellets were
added to the water and stirred continuously; the solid KOH appeared to dissolve and the
temperature readings rose​.​ The temperature was recorded for another three minutes, making a
total of five minutes recording. After the five minutes, this solution was poured into a clean, dry
beaker and covered with a glass watch glass in order to prevent contamination. The styrofoam
cup was then rinsed out with deionized water, three times to prepare for the next reaction. The
data collected was exported to an excel file and saved.
Another two to three grams of KOH pellets were measured out and recorded, then put
into the fifty mL beaker which was also measured and recorded. A new Microlab experiment
was prepared, recording time and temperature just as before, and the thermistor was set up
suspended over the styrofoam cup again. Approximately 160 mL of 1.5 M HCl was placed into a
250 mL beaker. 100 mL of that solution was measured using a graduated cylinder and poured
into the styrofoam cup. The temperature of the HCl was recorded for two minutes, and, just as
before, the solid KOH was added stirring continuously. The temperature was recorded for an
additional three minutes, for a total of five minutes. Then, the pH of the solution was tested using
a piece of pH paper, which remained pink; this result was then recorded. The styrofoam cup was
then rinsed three times with deionized water again and prepared for the next reaction. The data
recorded was then exported to an excel file and saved.
Next, the graduated cylinder was rinsed with deionized water and fifty mL of the KOH
solution from the first part of the experiment was measured out and placed into a clean, dry fifty
mL beaker. The graduated cylinder was rinsed and dried again and fifty mL of the 1.5 M HCl
solution was measured and poured into the styrofoam cup. The thermistor was suspended in the
fifty mL beaker containing the KOH solution for two minutes, and the initial temperature was
recorded. This data was exported and saved to an excel file. The thermistor was then rinsed and
dried and a new Microlab experiment was set up. The temperature of the HCl in the styrofoam
cup was then recorded for two minutes, and just as before, after two minutes, the KOH solution
was added and stirred continuously. The temperature was recorded for an additional three
minutes for a total of five minutes. The pH of this solution was also tested with a pH paper,
which turned blue in color. This data was exported and saved to an excel file, and the solution
was disposed of.
Data
The graphs below show the data collected in all three trials of the experiment. Graph 1
shows the time vs. temperature data collected for the first reaction with water and solid KOH.
Graph 2 shows the data collected in the second reaction of the lab with HCl and solid KOH.
Graph 3 shows the data collected for the final reaction of the lab, adding the KOH created in the
first reaction to HCl. Graph 4 shows the temperature of the KOH solution that was recorded on
its own for two minutes in the third part of the lab, before being added to the HCl solution. Table
1 shows various numerical data that was either collected or calculated, that gives information
such as the heat that the reaction gives off, or the limiting reagent in the reaction. Table 2 shows
the same data as Table 1 shared from the three other groups in the lab.
Graph 1
Graph 1: Shows the data collected from the thermistor for the first reaction of the experiment,
adding solid KOH to water. The temperature is represented in Celsius on the y-axis while the
time is on the x-axis in seconds. At 120 seconds, the KOH was added, resulting in a rise in
temperature, from roughly 23 degrees Celsius to about 26. A trendline was created to best fit the
data, which has a fairly low R-squared value of 0.7588.
Graph 2
Graph 2: Shows the data collected from the thermistor for the second reaction performed, adding
solid KOH to 1.5 M HCl solution. The temperature is represented in Celsius on the y-axis while
the time is on the x-axis in seconds. At 120 seconds, the KOH was added, resulting in a rise in
temperature, from roughly 24 degrees Celsius to about 33. A trendline was created to best fit the
data, which has a fairly low R-squared value of 0.7343.
Graph 3
Graph 3: Shows the data collected from the thermistor for the third reaction performed, adding
the KOH solution from part 1 to 1.5 M HCl solution. The temperature is represented in Celsius
on the y-axis while the time is on the x-axis in seconds. At 120 seconds, the KOH solution was
added, resulting in a rise in temperature, from roughly 23 degrees Celsius to about 27. A
trendline was created to best fit the data, which has a fairly low R-squared value of 0.7116.
Graph 4
Graph 4: Shows the temperature of the KOH solution created in part 1 that was collected over
two minutes. Just as the other graphs, the temperature is on the y-axis in Celsius, and the time in
seconds on the x-axis in The temperature stayed fairly consistent, fluctuating around 25 degrees
Celsius.
Table 1
Part 1
mass of KOH
moles of KOH
mL solution
ΔT
Heat(J)
Heat (J)/mole
KOH
Part 2
Part 3
2.234 g mass of KOH
2.541 mass of KOH
0.0398 mol moles of KOH
4.46
0.0453 moles of KOH
0.0769
0.15 moles of HCl
102.234 mL moles of HCl
3.157 ℃ mL of solution
0.075
102.541 mL of solution
1320.8 pH test
100
acidic pH test
33185.9 Limiting Reagent
basic
KOH Limiting Reagent
HCl
8.705 ΔT
ΔT
Heat (J)
3.32
98.426 Heat (J)
Heat(J)/ mole of
KOH
18.77
Heat(J)/ mole of
2127.75 KOH
235.8
Percent error=
1438%
Table 1: Data collected from all three trials that included the mass of the KOH, the calculated
moles of the KOH, the mL of solution, the change in temperature, the calculated heat produced,
and the calculated heat per mole produced. Part 2 and 3’s columns are longer because they
include the calculated moles of HCl, pH test, and the calculated limiting reagents.
Table 2
Group 1
Part 1
Part 2
Part 3
Mass of KOH
3.0684 Mass of KOH
2.9217 Mass of KOH
3.1111
Moles of KOH
0.0538 Moles of KOH
0.0513 Moles of KOH
0.0555
103.0684 Moles of HCL
0.15 Moles of HCL
0.0075
mL of solution
ΔT
Heat (J)
Heat (J)/ Mole of
KOH
4.6 mL of solution
1924.64 pH test
102.9217 mL of solution
Acidic
35773.98 Limiting Reagent KOH
ΔT
Heat (J)
Heat (J)/ Mole of
KOH
100
pH test
Basic
Limiting Reagent
HCL
10.6 ΔT
4435.04 Heat (J)
Heat (J)/ Mole of
86453.02 KOH
3.6
1506.24
27139.46
Percent error=
27.23%
Group 2
Part 1
Part 2
Part 3
Mass of KOH
2.75 Mass of KOH
2.32 Mass of KOH
Moles of KOH
0.049 Moles of KOH
0.041 Moles of KOH
0.098
0.15 Moles of HCl
0.075
mL of solution
ΔT
Heat (J)
Heat (J)/mole
KOH
102.75 Moles of HCl
mL of final
102.32 solution
3.71 mL of solution
1550.4 pH test
31632.3 Limiting reagent
5.5
100
acidic
pH test
basic
KOH
Limiting Reagent
HCl
7.84 ΔT
ΔT
Heat (J)
2.93
3279.3 Heat (J)
Heat (J)/mole
KOH
612.7
79307.1 Heat (J)/mole KOH
6252
Percent error=
52.2 %
Group 3
Part 1
mass of KOH
Part 2
2.5 mass of KOH
moles of KOH
0.0449 moles of KOH
ml of solution
100.67 Moles of HCl
ΔT
heat(J)
Heat(J)/mol KOH
Part 3
2.71 Mass of KOH
0.0483 Moles of KOH
33360.19 Limiting reagent
ΔT
Heat (J)
Heat (J)/mole
KOH
0.0898
0.15 Moles of HCl
0.075
mL of final
101.51 solution
3.58 mL of solution
1497.87 pH test
5.04
101.53
acidic
pH test
basic
KOH
Limiting Reagent
HCl
10.18 ΔT
4260.57 Heat (J)
88210.501 Heat (J)/mole KOH
0.351
146.858
1635.394
Percent error=
60.32%
Table 2: Shows the same data as displayed in Table 1 for three additional groups who performed
the lab.
Data Analysis and Calculations
mass of KOH(g)
Moles of KOH= M olar M ass of KOH(g/mol)
2.541g
56.108g/mol
= ​0.0453 moles of KOH
(1)
(volume of liquid (mL)) + (mass of solid(g)) = v​ olume of solution (mL)
(​100​mL H​2​O) + (​2.541​g KOH) = 1​ 02.541 ​mL of solution
(2)
(final temp.) – (initial temp.) = ​change in temp.​℃
(​32.965​℃) – (​24.26​℃) = 8​ .705​℃
(3)
Molarity of HCl ・Volume(L)=
​ Moles of HCl
1.5​M ・​0.1​L= ​0.15 moles of HCl
(4)
(moles of KOH)(mole ratio of KOH:H2​ O
from equation)(molar mass of H​2O)=
​ mass of H​2​O
​
​
mol H2O
(0.0453​mol KOH)​ ( 11 mol
)(18.016​g/mol)​ = 0.816 grams of H​2​O
KOH
(5)
(mass of H​2​O)(specific heat(constant))(change in temperature(℃))= H
​ eat(J)
(2.7024​g H​2O
​ (4.184)(8.705​℃)= 9​ 8.426​J
​ )
(6)
Heat(J)
mole KOH = Enthalpy
98.426J
= 2172.76 ​J/mol
0.0453mol KOH
(7)
Theoretically: ​(Enthalpy of reaction 1) + (Enthalpy of reaction 3)= ​Enthalpy of reaction 2
(33185.9​J/mol​) + (235.8​J/mol) ≠
​ 2172.76​J/mol
Experimentally: (33185.9​J/mol) + ​(235.8​J/mol​) = 33421.7​J/mol
​
(8)
|experimental value−theoretical value|
× 100 = % error
theoretical value
|(33421.7J/mol)−(2172.76J/mol)|
× 100 = 1438 % error
2172.76J/mol
(9)
The temperature change that was recorded for each of the reactions shows that each
reaction rose in temperature, indicating that the reaction gave off heat that was then absorbed and
registered by the thermistor, making them all, exothermic reactions. The graphs of the reactions
all resemble each other in appearance: a fairly constant line at one temperature, then an almost
vertical rise in temperature at the two minute mark, when the KOH was added, before abruptly
flattening off again for the remaining two minutes. However, they differ in both their initial
temperatures, as well as their magnitude of change. For example, the first reaction began at a
steady temperature of about 23 degrees Celsius, and rose to a temperature of about 26 degrees
Celsius. Whereas the second reaction began at a temperature around 24 degrees Celsius and rose
to approximately 33 degrees Celsius; becoming the experiment’s largest difference. Finally, the
third reaction began, again, around 23 degrees and rose to about 27 degrees Celsius.
When compared to the other groups data, the amount of KOH used, correlated to the
amount of heat derived from the reaction. Because our results were so vastly different from the
other group’s data, these approximations were calculated only using their data. It appears that for
the first two reactions, the higher mass of KOH used, resulted in a greater amount of heat being
released from the reaction; with the average ratio of grams of KOH to heat released, being
approximately 1.7g/1000J for the first reaction and about 0.67g/1000J. In regard to the third
reaction, the mass of KOH seemed to have no significant effect on the amount of heat released,
as the ratios differ from 2g/1000J in the first group, and 9g/1000J in the second group, to
34g/1000J in the third group.
Theoretically, based on Hess’s Law, the second reaction’s enthalpy should be the sum of
the first and third reaction’s enthalpy. This law seems proven by looking at the data from the
other three groups, all their percent error values coming in at under 100 percent. However, when
looking at the experimental data from the experiment we performed, Hess’s Law cannot be
accepted from this experiment with a percent error as high 1438 percent.
The acidity test that was performed and recorded for the second and third reactions was
significant to defining which reactant was the limiting reagent, confirming the values that were
then calculated using the amount of KOH used. The limiting reagent for the second reaction was
the KOH because the solution was acidic, meaning there was more of the acidic HCl in the
solution; and the limiting reagent for the third reaction was HCl because the solution was basic,
meaning there was more of the basic KOH in the solution.
Conclusion
The intention of this lab was to demonstrate the properties of temperature change in
chemical reactions and how that relates to heat and thermodynamics. These principles were
effectively conveyed using the thermistor and Microlab to record in a visually accessible manner
which made clear a significant temperature change in reactions that, visually, don’t make the
most dramatic transformations. Hess’s Law was applied using the enthalpies, calculated from the
heat released in the reaction. While the concept of Hess’s law was demonstrated in terms of
theoretical values in this lab, the calculated results did not support the law, yielding a 1438
percent error.
This extremely high value is likely due to a number of things that could have happened
during the experiment. For example, the difference in initial mass of KOH could have caused a
discrepancy in the heat release values, as was mathematically discovered, the amount of KOH
has a relationship with the heat released that is not necessarily always consistent. Additionally, a
reaction could have lost more heat than was recorded by the thermistor, particularly in reaction
two, where the calculated heat released made the least sense. It is also likely that the KOH did
not completely dissolve in the solution, meaning that the full chemical reaction did not occur,
making the heat that was produced a result of something other than the full intended reaction.
And although a watch glass was placed over the aqueous KOH in order to prevent
contamination, it is possible that that solution became contaminated while sitting, which would
alter the results of the third reaction.
The heat values collected for each part of the experiment went in descending order, the
first reaction producing the most heat, and the third producing the least. However, based on the
chemical equations of the reactions, part two produces the same products as the solution made in
part one, combined with the HCl; meaning that part two is essentially a combination of parts one
and three. Because of this chemical relationship, Hess’ Law states that the heat values of parts
one and three should add up to be the heat value of part two. Meaning that the descending heat
values from one to three would mathematically make this impossible, proving, as was stated
above, that Hess’ Law is not supported by the data collected.
References
Cohen, S. (2019, September 30). Hess’s Law. Retrieved from
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook
_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Thermodynamics
/Thermodynamic_Cycles/Hess’s_Law.
Harshini, M., & Acacia, C. (2015, December 3). Enthalpy of Reaction and Hess’s Law. Retrieved
from https://sites.google.com/site/experiment6enthalpy/home.
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