Experiment 1

BASIC CALORIMETER EXPERIMENT: HEAT TRANSFER OF WATER

I. INTRODUCTION
Let's delve into the fascinating world of heat and temperature, and how they play a crucial role
in our everyday experiences. Imagine you're sipping a comforting cup of hot cocoa in a foam
cup. You know the cocoa warms up, but the cup remains cool to the touch. Ever wondered why?
It's all about the way heat is transferred, and it's a concept that becomes even more intriguing
when we venture into calorimetry.
Calorimetry, the science of measuring heat transfer, often involves insulated cups, like the ones
used in experiments. These special cups are designed to minimize heat exchange with the
surroundings, ensuring precision in the said experiment. But before we dive into the world of
calorimetry, let's establish a solid foundation by understanding the distinctions between heat and
temperature.
Heat isn't quite the same as temperature. Temperature tells us how hot or cold something is,
while heat represents the total energy contained within a substance. It's a result of multiplying
temperature, mass, and the specific heat of the material. When substances with different
temperatures mix, heat energy is exchanged, and the rate of this exchange depends on the
mass and specific heat of each material.
This laboratory activity is conducted by other groups, and we will only be analyzing their data
and determining the contrasting results they had with ours, enabling us to come up with a more
efficient and profound understanding of the topics that are expected of us to successfully turn
from theoretical to practical applications. So, let's embark on this journey to demystify heat,
temperature, and calorimetry, exploring the subtle intricacies that govern our experiences with
warmth and cold, and how we can harness these principles for accurate and insightful
experiments.
II. OBJECTIVES
1. To conduct basic calorimeter experiment.
2. To apply the basic knowledge in solving heat energy.
3. To compare the heat energy of cold and hot water.
III. EQUIPMENT AND SUPPLIES
 Styrofoam cup
 cold and hot water
 thermometer
 balance
IV. PROCEDURE
1. Measure the mass of the empty calorimeter with a balance. Record on a data table.
2. Pour cold water ---no ice—into the calorimeter until it is one third full. Find the total mass
of the calorimeter and cold water. Record the mass on the data table.
3. Put the lid on the calorimeter and push a thermometer through the slit in the lid. Make
sure the thermometer reaches the water. Measure the temperature of the cold water.
4. Repeat steps one, two and three, this time using hot water. The hot water should be a
minimum of 50 degrees Celsius.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 5. Pour the hot water from the Calorimeter into the cold water in the second calorimeter.
Quickly close the lid to reduce unwanted heat loss.
6. Push the thermometer through the hole in the lid and observe the temperature of the
mixed water.
7. Once the temperature stops changing, record in data chart.
8. Repeat the experiment twice with different masses of water.
Experiment 1
Basic Calorimeter Experiment: Heat Transfer of Water
REPORT SHEET
V. DATA AND OBSERVATION
Table 1.1 Heat Transfer of Water
GROUP 1
MASS (g) TEMPERATURE (℃ )_
Trial 1 Trial 2 Trial 1 Trial 2
Empty Calorimeter 4.23 g 4.50 g 26 ℃ 23 ℃
Calorimeter and Cold 84.44 g 79.42 g 4℃ 4℃
Water
Calorimeter and Hot 87.48 g 78.26 g 52 ℃ 52 ℃
Water
Calorimeter with Hot and 169.01 g 152.56 g 14 ℃ 28 ℃
Cold Water

Calculations and Review Questions:

1. Complete calculations to find the total mass of hot plus cold water. Calculate the
temperature change of cold water after mixing. Repeat with hot water data.
Total mass = 87.48 g + 84.44 g = 171.92 g
Temperature Change of Cold Water = 4°C - 4°C = 0°C
Temperature Change of Hot Water = 52°C - 52°C = 0°C
2. Using this information, calculate the heat energy of cold water by using the following
equations: Heat energy of cold water equals the mass of cold multiplied by temperature
change of cold water, multiplied by the specific heat of water which is one calorie per
gram.
Q=mC ∆ T
mcoldwater =84.44 g

cal
C=1
g℃
∆ T coldwater =0 ℃

cal
Qcoldwater =(84.44 g)(1 )(0 ℃)
g℃

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Qcoldwater =0 cal

3. Repeat with hot water using the hot water data to determine its final heat energy.
Q=mC ∆ T
mhot water =87.48 g

cal
C=1
g℃
∆ T hotwater =0 ℃

cal
Qhotwater =(87.48 g)(1 )(0 ℃)
g℃
Qhotwater =0 cal
4. Find also the heat energy of the mixed water.
Q=mC ∆ T
mmixed =171.92 g

cal
C=1
g℃
∆ T mixed =28 ℃−14 ℃=14 ℃

cal
Qmixed =(171.92 g)(1 )(14 ℃)
g℃
Qmixed =2406.88 cal

VI. CONCLUSION AND RECOMMENDATION

The data collected from the calorimeter experiment provides valuable insights into heat transfer
between substances. Several key observations can be made from the experiment:
1. Heat Absorption and Release: The experiment demonstrates the principles of heat
absorption and release. When hot water was added to the calorimeter with cold water,
there was a transfer of heat from the hot water to the cold water, resulting in a
temperature change in both.
2. Temperature Equilibration: The final temperature of the mixed hot and cold water in
the calorimeter was not simply the average of the initial temperatures. Instead, it reflects
the equilibration of heat between the two water sources. This indicates that heat transfer
depends on both the quantity of material and its initial temperature.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 3. Calorimeter Efficiency: The experiment also highlights the importance of an insulated
calorimeter. The calorimeter effectively minimized heat loss to the surroundings, allowing
for a more accurate measurement of heat transfer.
Based on the observations and results of the calorimeter experiment, the following
recommendations can be made:
1. Calorimeter Calibration: It is important to ensure the calorimeter is properly calibrated
before conducting experiments. Accurate measurements of mass and temperature are
crucial for precise heat transfer calculations.
2. Multiple Trials: Conducting multiple trials of the experiment can help reduce the impact
of random errors and improve the reliability of the data. Averaging the results from
multiple trials can provide more accurate values for heat transfer.
3. Controlled Conditions: Maintain controlled conditions during the experiment, including
the use of a consistent water source and the elimination of external heat sources that
may affect the results.
4. Thorough Mixing: Ensure thorough mixing of hot and cold water in the calorimeter to
facilitate heat equilibration and accurate temperature measurements.
5. Calorimeter Insulation: Continue to use well-insulated calorimeters to minimize heat
loss to the surroundings, as this is crucial for accurate heat transfer experiments.
In summary, the calorimeter experiment demonstrates the principles of heat transfer and the
importance of controlled conditions and accurate measurements. By following the
recommendations, future experiments can yield even more precise results, contributing to a
better understanding of heat transfer processes.

REFERENCES:
1. Clark, J. (n.d.). Calorimetry. Chemguide. Retrieved from
http://www.chemguide.co.uk/physical/energetics/calorimetry.html
2. Carnegie Mellon University. (n.d.). Virtual Labs - Calorimetry. ChemCollective. Retrieved
from http://chemcollective.org/virtuallabs.php
3. Khan Academy. (n.d.). Calorimetry. Retrieved from
https://www.khanacademy.org/science/chemistry/thermodynamics-chemistry/enthalpy/
v/calorimetry

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Experiment 2
HEAT OF SOLUTION
I. INTRODUCTION
Picture this scenario in your everyday life where you stir sugar into your morning coffee or tea.
You've likely noticed how this simple act results in a change in the beverage's temperature.
What you may observe is the intriguing phenomenon known as the "heat of solution," or in more
scientific terms, the "enthalpy change of solution."
This captivating concept revolves around the heat, either absorbed or released, when a solute,
such as sugar, is introduced into a solvent, like water. It is a dynamic interplay of energy, which
occurs during the dissolution process, which can be complex.
Let's dive into the intricacies of this fascinating process:
1. Absorption of Energy (Endothermic): As you start dissolving the sugar, energy is absorbed
into the system. This part of the process is known as "endothermic." Think of it as the energy
required to break the bonds holding the solute molecules (in this case, sugar) together and the
bonds between the solvent molecules (the water molecules).
2. Release of Energy (Exothermic): As the solute particles intermingle with the solvent
particles, they form new bonds. This scenario illustrates an "exothermic" process, where energy
is released. The creation of these fresh bonds between the solute (sugar) and the solvent
(water) leads to the release of energy.
The heat of solution, represented as Δ H solution, quantifies the net difference between the energy
absorbed and the energy released per mole of solute during this process. It tells us whether the
overall process is endothermic (positive Δ H solution,) or exothermic (negative Δ H solution,).
- Positive Heat of Solution (Endothermic): Some solutes demand more energy to break their
existing bonds than they release when forming new bonds with the solvent. This means that
dissolving these solutes absorbs more energy than it releases, causing a decrease in the
solution's temperature. So, adding such solutes cools down the solution.
- Negative Heat of Solution (Exothermic): In contrast, other solutes require less energy to
break their existing bonds and release more energy when forming new bonds with the solvent.
Dissolving these solutes releases more energy than it absorbs, leading to a rise in the solution's
temperature. In this case, adding the solute warms up the solution.
To precisely measure the heat of the solution, scientists use units like kilojoules per mole
(kJ/mol). While older sources might use units like kilocalories per mole (kcal/mol or C/mol),
where 1 kilojoule is roughly equivalent to 0.239 kilocalories, and 1 kilocalorie is approximately
4.1841 kilojoules.
In experimental settings, chemists determine the heat of solution by dissolving a known amount
of a solute in a known amount of a solvent and meticulously measuring the resulting

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
temperature change. This approach allows them to unveil the intricate energy dynamics at play
during the dissolution process, helping us understand how substances interact and how these
interactions can either heat up or cool down a solution.
This experiment was conducted by Groups 1 and 4. We will be analyzing the results of their
experiments, especially the data they have gathered, and compare them with each other. By
doing so, we could expect that the theoretical knowledge is able to be applied efficiently and
allow for a profound understanding of the concept, aiding us to arrive at an excellent and
accurate conclusion of our own.
II. OBJECTIVES
1. To measure the heat of solution of some substances.
2. To compare the experimental and theoretical values of heat of solution.
III. EQUIPMENT AND SUPPLIES
 Goggles
 Gloves
 Protective Clothing (Lab Gowns)
 Balance and Weighing Papers
 Calorimeter
 Thermometer
 Graduated Cylinder
 100 mL ammonium nitrate (40.0 g = 0.5 mol) – Due to the unavailability of the material,
they used Magnesium Sulfate as an alternative
 Sodium Chloride (29.2 g = 0.5 mol)
 Sodium Hydroxide (20.0 g = 0.5 mol)
 Water (at room temperature)
IV. PROCEDURE
In order to ensure the precision of their measurements, the solutes, the solvent (water), and the
calorimeter were carefully maintained at room temperature. They had thoughtfully replaced cold
tap water with water that had been allowed to equilibrate to room temperature within a 1.5-liter
container.
With their safety gear on, including splash goggles, gloves, and lab coats, our colleagues began
the experiment.
In a slight twist of fate, it turned out that ammonium nitrate, the intended solute, was
unavailable. In response, they made the decision to proceed with magnesium sulfate, as an
equally valid substitute. A meticulous measurement of precisely 40.0 grams of magnesium
sulfate was executed, with the mass duly recorded. With their substitute in place, they
proceeded with their experiment.
Using a graduated cylinder, they measured out a precise 100.0 milliliters (equivalent to 100.0
grams) of water. The carefully measured water was then introduced into the calorimeter, and
they ensured that the cover was securely replaced to maintain the integrity of their
measurements.
With the calorimeter primed and the temperature of the water recorded, they were all set to kick
off their calculations with precision.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
The stage was set for the addition of 40.0 grams of magnesium sulfate into the calorimeter.
Their actions were swift, and they ensured the calorimeter was securely sealed.
Following this, they carried out a thorough stirring of the solution, or alternately, gently swirled
the calorimeter to aid the complete dissolution of the magnesium sulfate. Their approach was
characterized by a methodical and detail-oriented execution.
As they watched the thermometer, anticipating the peak temperature change, they were
captivated by the heat transfer process at play. Once the temperature change reached its
zenith, they diligently recorded this pivotal data point. The temperature difference, a critical
aspect of their experiment, was meticulously calculated and duly noted.
With the experiment involving magnesium sulfate successfully completed, they turned their
attention to the safe disposal of the solution and ensured that the calorimeter was meticulously
rinsed, ready for the next steps.
In their commitment to rigor and precision, they replicated the experiment following the same
protocol. For this iteration, they substituted the magnesium sulfate with 29.2 grams of sodium
chloride, adhering to the specified requirements. Their approach remained unwavering, marked
by exact measurements and careful observation.
In the final leg of their experiment, they replicated the procedure yet again. This time, they
substituted the magnesium sulfate with 20.0 grams of sodium hydroxide, maintaining their
meticulous approach throughout. Their diligence and precision in data recording were consistent
with their previous efforts.
As our colleagues concluded their experiment, they took into consideration the responsible
disposal of their solutions. According to laboratory guidelines, all solutions could be safely
flushed down the drain with an ample supply of water. In this manner, their experiment reached
its conclusion, providing them with valuable insights into heat transfer and the intricacies of
enthalpy changes.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Experiment 2
Basic Calorimeter Experiment: Heat Transfer of Water
REPORT SHEET
V. DATA AND OBSERVATION
Table 1.1 Heat Transfer of Water
GROUP 1
Solute Mass(g) Water Solution Tempera Calculated Heat
Temperature Temperature ture of Solution,
(℃ ) , (℃ ) Differenc kJ/mol
e (℃ )
A. Magnesium 112.09 g 28 ℃ 26℃ -2 ℃ kJ
−3
Sulfate
−8.368 ×10
mol
B. Sodium 118.66 g 28 ℃ 27℃ -1 ℃ −3 kJ
Chloride
−4.184 ×10
mol
C. Sodium 121.97 g 28 ℃ 96 ℃ 68 ℃ kJ
Hydroxide
0.284512
mol

GROUP 4
Solute Mass(g) Water Solution Temperatur Calculated Heat of
Temperature Temperatur e Solution, kJ/mol
(℃ ) e (℃ ) Difference
(℃ )
A. Magnesiu 10.03 g 30 ℃ 30 ℃ 0℃ kJ
m Sulfate
0
mol
B. Sodium 10.01 g 30 ℃ 26 ℃ −4 ℃ kJ
Chloride
−0.016736
mol
C. Sodium 10 g 30 ℃ 49 ℃ 19 ℃ kJ
Hydroxide
0.079496
mol

Calculations and Review Questions:

1. Using your experimental data, calculate the heats of solution for ammonium nitrate,
sodium chloride, and sodium hydroxide. Enter your calculated values in Table 2-1.
SOLUTE GROUP 1 GROUP 4
A. Magnesium Sulfate Q=mC ∆ T Q=mC ∆ T
Q= ( 40.0 g ) ¿ procedure Q= ( 40.0 g ) ¿ procedure
J J
C=4.184 C=4.184
g℃ g℃
∆ T =−2 ℃ ∆ T =0℃
Q= ( 40.0 g ) ¿ Q= ( 40.0 g ) ¿
¿(−2℃) ¿(0℃ )

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 −334.72 J 0J
40.0 g 40.0 g
Δ H solution= Δ H solution=
1000 1000
−3 kJ kJ
Δ H solution=−8.368 × 10 Δ H solution=0
mol mol
B. Sodium Chloride Q=mC ∆ T Q=mC ∆ T
Q= ( 29.2 g ) ¿ procedure Q= ( 29.2 g ) ¿ procedure
J J
C=4.184 C=4.184
g℃ g℃
∆ T =−1 ℃ ∆ T =−4 ℃
Q= ( 29.2 g ) ¿ Q= ( 29.2 g ) ¿
¿(−1℃) ¿(−4 ℃)
−122.1728 J −488.6912 J
29.2 g 29.2 g
Δ H solution= Δ H solution=
1000 1000
−3 kJ kJ
Δ H solution=−4.184 ×10 Δ H solution=−0.016736
mol mol
C. Sodium Hydroxide Q=mC ∆ T Q=mC ∆ T
Q= ( 20.0 g ) ¿ procedure Q= ( 20.0 g ) ¿ procedure
J J
C=4.184 C=4.184
g℃ g℃
∆ T =−1 ℃ ∆ T =19 ℃
Q= ( 20.0 g ) ¿ Q= ( 20.0 g ) ¿
¿(68℃) ¿(19℃)
5,690.24 J 1,589.92 J
20.0 g 20.0 g
Δ H solution= Δ H solution=
1000 1000
kJ kJ
Δ H solution=0.284512 Δ H solution=0.079496
mol mol

2. Look up published values for the heats of solution of those three compounds on the
Internet or in a printed reference. How closely do the values that you obtained
experimentally correspond to the published values? If your values are significantly
different, propose possible explanations.

 For Magnesium Sulfate (MgSO4):

−3 kJ
Experimentally Obtained (Group 1): Δ H solution=−8.368 × 10
mol
kJ
Experimentally Obtained (Group 4): Δ H solution=0
mol
Published Value: -91.2 kJ/mol

The experimentally obtained values for both Group 1 and Group 4 are significantly
different from the published value of -91.2 kJ/mol. The discrepancy may result from

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 various factors, including variations in experimental conditions, impurities in the
magnesium sulfate used, or equipment calibration.

 For Sodium Chloride (NaCl):

−3 kJ
Experimentally Obtained (Group 1): ΔH = Δ H solution=−4.184 ×10
mol
kJ
Experimentally Obtained (Group 4): Δ H solution=−0.016736 l
mol
Published Value: 3.9 kJ/mol

The experimentally obtained values in both Group 1 and Group 4 are significantly
different from the published value of 3.9 kJ/mol. The variation could be attributed to
factors such as equipment calibration, the purity of sodium chloride, or differences in
the experimental setup.

 For Sodium Hydroxide (NaOH):

kJ
Experimentally Obtained (Group 1): Δ H solution=0.284512
mol
kJ
Experimentally Obtained (Group 4): ΔH = Δ H solution=0.079496
mol
Published Value: -44.5 kJ/mol

The experimentally obtained value in Group 1 is significantly different from the

published value of -44.5 kJ/mol. Possible reasons for this difference include
variations in experimental conditions, equipment calibration, or the purity of sodium
hydroxide.

In summary, the experimentally obtained values for the heat of solution differ significantly
from the published values for all three compounds. These variations may be due to
factors such as impurities, equipment calibration, procedural differences, or variations in
experimental conditions. To improve the accuracy and reliability of the results, it is
essential to critically examine and address potential sources of error in the experimental
setup. Additionally, conducting experiments under controlled and standardized
conditions is crucial for obtaining data that closely corresponds to the published values.
VI. CONCLUSION AND RECOMMENDATION
Upon analyzing the data from two different groups conducting basic calorimeter experiments to
measure the heat transfer of water, several key observations and conclusions can be drawn:
1. Variability in Results: It is evident from the data that there are significant differences
in the temperature changes and, consequently, the calculated heat of solution ( Δ H solution
,) for the same solutes in different groups. For example, the temperature difference for
the dissolution of sodium hydroxide in Group 1 was 68°C, while in Group 4, it was 19°C.
This discrepancy in results suggests that external factors or errors might have influenced
the outcomes.
2. Magnesium Sulfate: In both groups, magnesium sulfate demonstrated a temperature
change close to zero. This is consistent with the known properties of this substance,

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 which has a relatively low heat of solution. It is expected to absorb or release minimal
energy during dissolution.
3. Sodium Chloride: The dissolution of sodium chloride showed varying temperature
changes in the two groups. In Group 1, it resulted in a 1°C change, while in Group 4, it
caused a 4°C change. These differences could be attributed to experimental errors or
variations in the precise measurement of the solute.
4. Sodium Hydroxide: The most significant variation is observed in the dissolution of
sodium hydroxide. In Group 1, the temperature change was 68°C, while in Group 4, it
was 19°C. Such a notable difference could be due to variations in the concentration of
the sodium hydroxide solutions used in the two groups, as higher concentrations tend to
produce more pronounced temperature changes.
From the conclusion that I was able to formulate, here are a few recommendations that I can
propose regarding the said experiment:
1. Calibration and Standardization: It's crucial to ensure that the equipment used in
different groups is properly calibrated and standardized. Variations in the calorimeter's
performance can introduce discrepancies in results. By establishing and following a
calibration protocol, more consistent and reliable data can be obtained.
2. Standardized Concentrations: To minimize the impact of variations in solute
concentrations, it is recommended to use standardized solutions. This ensures that the
same concentration is used across different groups, leading to more comparable results.
3. Measurement Precision: To reduce errors in measurements, ensure that the mass of
solute and solvent is measured with high precision. Small measurement errors can
significantly affect the calculated heat of solution.
4. Data Verification: Consider implementing a process for data verification and validation
to catch and correct any errors or inconsistencies in the recorded data.
5. Replication: Conducting experiments in triplicate or more can help to detect outliers and
improve the reliability of the results. This can also reveal whether variations in the
experiments are consistent or sporadic.
In summary, the significant differences in results between Group 1 and Group 4 are likely due to
factors such as calibration, concentration, and measurement precision. By implementing the
recommended measures, more consistent and accurate data can be obtained in future
calorimeter experiments.

REFERENCES:
1. Libretexts. (2022, August 8). 17.13: Heat of Solution. Chemistry LibreTexts.
https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Introductory_Chemistry_
(CK-12)/17%3A_Thermochemistry/17.13%3A_Heat_of_Solution
2. Stubbings, J. (n.d.). Heat of Solution Chemistry tutorial.
https://www.ausetute.com.au/heatsolution.html
3. Libretexts. (2021, August 23). Chapter 9.5: Enthalpies of Solution. Chemistry LibreTexts.
https://chem.libretexts.org/Courses/Howard_University/General_Chemistry
%3A_An_Atoms_First_Approach/Unit_4%3A__Thermochemistry/
09%3A_Thermochemistry/Chapter_9.05%3A_Enthalpies_of_Solution

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Experiment 3
HEAT OF FUSION OF ICE
I. INTRODUCTION
In the realm of thermodynamics, the behavior of substances at their phase transitions is a
subject of profound significance. When heat is introduced into a system comprising various
phases of a substance, particularly at the melting point of a solid phase, a peculiar phenomenon
emerges. Contrary to the intuitive expectation of an immediate rise in temperature, the system's
temperature remains constant. This is an intriguing manifestation of energy exchange,
characterized by the complete liquefaction of the solid substance.
This distinctive occurrence transpires because the introduced heat energy is primarily absorbed,
without inducing a simultaneous temperature increase. It is absorbed to an extent sufficient to
weaken the intermolecular attractions holding the atoms or molecules within the solid lattice.
This absorption of energy without a corresponding change in temperature is a fundamental
concept in thermodynamics known as the heat of fusion. This property may also be referred to
as the enthalpy of fusion or specific melting heat.
The precise quantification of the heat of fusion is an essential endeavor, and it finds particular
relevance in the discipline of physical chemistry. It is a parameter that distinguishes various
substances and plays a pivotal role in the study of phase transitions, particularly in solid-to-liquid
transformations. Our laboratory activity focuses on the meticulous measurement of the heat of
fusion, a process of great scientific import.
Notably, the substance involved in this experiment is ice, which is composed of hydrogen
atoms. Ice is distinguished by its notable heat of fusion, primarily attributable to the substantial
intermolecular attractions inherent in the hydrogen bonding network.
In scientific notation, the unit of measurement for the heat of fusion is 'Joules per mole' (J/mol)
within the International System of Units (SI). Alternative units such as 'calories per gram' (cal/g),
'kilojoules per kilogram' (kJ/kg), and even 'British Thermal Units per pound' (BTU/lb) are
available. For the purposes of this experiment, we will adhere to the widely recognized unit
'cal/g,' a format still commonly employed in contemporary scientific discourse.
II. OBJECTIVES
1. To efficiently measure the heat of fusion of ice.
2. To conduct a basic application of theoretical knowledge regarding the heat of fusion.
III. EQUIPMENT AND SUPPLIES
For this experiment, our group made use of the following supplies and laboratory apparatuses:
 Goggles
 Gloves
 Protective clothing
 Calorimeter (Considering the availability of materials, we used a coffee-cup calorimeter)
 Thermometer
 Graduated Cylinder
 100 ml ice bucket or similar wide-mouth container beaker
 600 ml beaker (In our case, only 500 ml beaker is available)
 Tongs

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
  Hotplate
 Ice (Crushed)
 Hot tap water
IV. PROCEDURE
Ensuring laboratory safety is paramount in any experimental work, particularly when handling
fragile or high-temperature equipment and working with potentially hazardous chemicals. To
address safety concerns, we commenced our experiment by donning appropriate protective
gear, including splash goggles, gloves, and laboratory gowns. Additionally, we organized our
workspace meticulously to prevent any accidental mishaps. With safety as our foremost priority
and materials properly arranged, we proceeded with the experiment.
Following the prescribed procedure, we initiated the experiment by taking a 500 mL beaker and
filling it approximately halfway with the previously prepared crushed ice. Subsequently, we
allowed the ice to transition into its liquid form. In parallel, a second 500 mL beaker was filled
with 400 mL of tap water. The latter was then placed on a hotplate and heated to a controlled
temperature of 65 ℃ .
Upon achieving the desired temperature in the water-filled beaker, we carefully poured
approximately 100 mL of the heated water into a 1000 mL graduated cylinder, utilizing a cloth
due to the unavailability of beaker tongs. To ensure uniformity and as stated in the directions,
we employed a 30-second timer before disposing of the water. This step was executed twice,
solely to preheat the graduated cylinder. Concurrently, we separated the solid ice from the ice-
water slush mixture and loaded the calorimeter to approximately half its capacity with ice. The
calorimeter was promptly sealed with its lid to prevent heat exchange with the surrounding
environment.
Subsequently, we introduced approximately 25 mL of the heated water into the graduated
cylinder. Following this addition, we meticulously measured the water's volume to the nearest
0.1 mL, recording this value in line A of the data table. The temperature of the water in the
graduated cylinder, designated as the initial temperature, was equally measured and recorded
to the nearest 0.1 ℃ on line B.
After documenting the initial volume and temperature, we took precautions to decant any
excess water from the calorimeter, and subsequently added 25 mL of the warm water from the
graduated cylinder to the calorimeter. The calorimeter was securely covered, and the mixture
was gently stirred. Thereafter, the temperature of the resulting ice-water mixture was precisely
measured and recorded to the nearest 0.1 ℃ on line C, denoted as the final temperature. After
this measurement, the liquid contents of the calorimeter were carefully transferred back into the
graduated cylinder, and the final volume was recorded on line D.
Following the first trial, we proceeded to conduct the second and final trial, replicating all
procedures exactly as previously performed. Although our original plan included up to four trials,
time constraints necessitated that we limit our experiment to a minimum of two trials.
As a final step, strict hygiene protocols were adhered to. To prevent any contamination, we
ensured that all laboratory equipment, as well as the laboratory workspace itself, was thoroughly
cleaned and restored to its original state.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Experiment 3
Heat of Fusion of Ice
REPORT SHEET
V. DATA AND OBSERVATION
Table 3.1: Heat of Fusion of Ice – observed and calculated data.
Item Trial 1 Trial 2
A. Volume, Initial, Ml 25.1 mL 25.9 mL
B. Temperature, initial, ℃ 50.1 ℃ 53.0 ℃
C. Temperature, final, ℃ 1.1 ℃ 2.0 ℃
D. Volume, final, Ml 50.1 mL 58.9 mL
E. Temperature Change (B-C), ℃ 49 ℃ 51 ℃
F. Volume Change (D-A), Ml 25 mL 33 mL
G. Heat of Fusion of Ice, cal/g cal cal
ΔHf =49 ΔHf =51
g g

Review Questions and Calculations

1. Using your experimental data for each of the first two trials, calculate the heat of fusion
of ice in cal/g and record your calculated values on line G of Table 3.1. The actual value
for the heat of fusion of ice is 79.72 cal/g. If the value you obtained experimentally differs
significantly, propose possible reasons for this variation.
Trial 1 Trial 2
m :25 mL=25 grams m 33 mL=33 grams
J cal J cal
Specific heat of water : 4.184 =1 Specific heat of water : 4.184 =1
g∙℃ g℃ g∙℃ g℃
∆ T =49 ℃ ∆ T =51 ℃
q=m Cw ∆ T q=m Cw ∆ T
cal cal
q=(25 g)(1 )(49 ℃) q=(33 g)(1 )(51 ℃)
g℃ g℃
q=1,225 cal q=1,683 cal
1,225 cal 1,683 cal
ΔHf = ΔHf =
25 g 33 g
cal cal
ΔHf =49 ΔHf =51
g g

The value we obtained experimentally indeed differs significantly from the actual value. The
possible reasons for these variations could be the following:
1. Experimental Error: The presence of random errors originating from imprecisions in
temperature, volume, and mass measurements may introduce deviations in the
calculated heat of fusion values. These errors, stemming from inherent limitations in

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 measurement devices or techniques, can propagate through the calculations and result
in disparities from the expected value.
2. Calorimeter Inefficiency: Calorimeters, although designed to minimize heat exchange
with the external environment, may not achieve perfect insulation. Consequently, heat
loss to the surroundings is a plausible factor contributing to variations in experimental
outcomes. Such inefficiencies can affect the accuracy of temperature measurements
within the calorimeter, ultimately influencing the determination of the heat of fusion.
3. Impurities: The presence of impurities within the ice or water used in the experiment
introduces a further layer of complexity. These impurities may alter the thermodynamic
behavior of the system, impacting the phase transition from ice to water. The resultant
heat of fusion can diverge from the expected value due to the unique interactions and
energy requirements associated with impurities.
4. Stirring and Mixing: The homogeneity of the system and the effectiveness of mixing are
critical in achieving accurate temperature measurements. Inadequate stirring can lead to
temperature gradients within the calorimeter, affecting the precision of the calculated
heat of fusion. Variations in temperature due to localized concentrations of ice or water
may confound the experimental results.
5. Systematic Errors: Systematic errors may emerge from limitations within the equipment
or the experimental procedure itself. These errors, which are consistent and repeatable,
can stem from equipment calibration discrepancies or procedural imperfections. Such
systematic deviations from ideal conditions can manifest in inconsistent heat of fusion
values.
VI. CONCLUSION AND RECOMMENDATION
From the experimental activity conducted regarding the heat of fusion of ice, it is evident that the
precise measurement of various parameters significantly influences the accuracy of
computations, particularly when determining the heat of fusion. The key findings of this
experiment can be summarized as follows:
1. Measurement Precision: The results underscore the critical importance of precision in
measuring temperature, volume, and mass when calculating the heat of fusion. Small
errors in these measurements can introduce notable variations in the final calculated
values.
2. Calorimeter Efficiency: The experiment has also highlighted the relevance of
calorimeter efficiency. Imperfections in the calorimeter's insulation may lead to heat loss
to the surroundings, thereby impacting the accuracy of temperature measurements and,
subsequently, the heat of fusion determination.
3. Impurity Effects: The presence of impurities within the experimental substances, in this
case, ice and water, introduces an additional layer of complexity. Impurities can alter the
thermodynamic behavior of the system, leading to deviations in the heat of fusion.
4. Mixing and Stirring: Inhomogeneity within the system due to inadequate mixing and
stirring has been identified as a potential source of error. Localized variations in the
concentration of ice or water may result in inconsistent temperature readings.
In light of the experiment's outcomes, several recommendations can be made to enhance the
accuracy and reliability of future experiments:

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 1. Measurement Precision: To mitigate the impact of measurement errors, it is imperative
to employ calibrated instruments with the highest precision available. Careful
measurement techniques and the use of redundant measurements can help reduce
measurement-related discrepancies.
2. Calorimeter Calibration: Calorimeters should be regularly calibrated to ensure optimal
insulation and minimize heat loss to the surroundings. Routine maintenance and
insulation checks are essential to maintain accurate temperature measurements.
3. Impurity Control: When working with substances such as ice and water, it is advisable
to use highly pure reagents to minimize the influence of impurities. Conducting purity
checks and employing purified substances can yield more consistent results.
4. Homogeneous Mixing: Vigorous and thorough mixing of the substances within the
calorimeter is crucial to achieve homogeneity and uniform temperature distribution.
Adequate stirring ensures that temperature measurements are representative of the
entire system.
5. Multiple Trials: Conducting multiple trials of the experiment can help reduce the impact
of random errors. A larger dataset allows for the identification of trends and the
calculation of more reliable average values.
6. Data Analysis Techniques: Implement robust data analysis techniques, such as error
propagation, to assess and account for uncertainties in measurements and calculations.

REFERENCES:
1. Tro, N. J. (2016). Chemistry: A Molecular Approach. Pearson.
2. Skoog, D. A., West, D. M., Holler, F. J., & Crouch, S. R. (2013). Fundamentals of
Analytical Chemistry. Cengage Learning.
3. Petrucci, R. H., Herring, F. G., Madura, J. D., & Bissonnette, C. (2016). General
Chemistry: Principles and Modern Applications. Pearson.
4. Silberberg, M. S. (2016). Chemistry: The Molecular Nature of Matter and Change.
McGraw-Hill Education.
5. Atkins, P., de Paula, J., & Keeler, J. (2017). Atkins' Physical Chemistry. Oxford University
Press.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Experiment 4
SPECIFIC HEAT OF METAL
I. INTRODUCTION
The fundamental concept of heat transfer governs the behavior of materials at varying
temperatures. When substances of disparate thermal states come into contact, heat
spontaneously flows from the higher temperature material to the lower one until thermal
equilibrium is achieved, where both entities share an identical temperature. Within the confines
of a controlled environment, exemplified by our calibrated calorimeter, the guiding principle
asserts that heat exchange with the surrounding milieu is virtually negligible. While practical
conditions might introduce minor deviations, these perturbations remain inconsequential in
relation to the substantial thermal changes we intend to quantify in this laboratory experiment.
The focal point of our current laboratory investigation centers on the precise determination of
the specific heat capacities of two distinct metallic elements: lead and iron. However, due to the
unavailability of such materials, we used stainless-steel bearings as an alternative. This
scientific objective unfolds by subjecting a precisely quantified mass of each metal to controlled
heating until they reach known temperatures. Subsequently, the heated metal is submerged into
a precisely measured quantity of cooler water, also characterized by its initial temperature. The
system is then allowed to naturally progress towards thermal equilibrium. During this thermal
interaction, heat energy is transferred from the metal to the surrounding water, invariably leading
to an observable elevation in the water's temperature.
The overarching scientific principle governing this exploration is that the quantity of heat lost by
the metal corresponds to the heat gained by the water, thereby enabling us to determine the
specific heat capacity of the metal. This fundamental concept forms the cornerstone of our
experimental design, facilitating the quantitative assessment of the specific heat capacity of the
stainless-steel bearings.
II. OBJECTIVES
1. To determine the specific heat of metals.
2. To compare the experimental and theoretical values of the specific heat of metals
3. To analyze the experimental sources of error.
III. EQUIPMENT AND SUPPLIES
For this experiment, our group made use of the following supplies and laboratory apparatuses:
 Goggles
 Gloves
 Protective clothing
 Balance and Weighing Boat
 Calorimeter (Considering the availability of materials, we used a coffee-cup calorimeter)
 Thermometer
 Graduated Cylinder, 100 mL
 Beaker, 500 mL (Supposedly 600 mL, but unavailable)
 Hotplate
 Stainless Steel Bearings (Since Lead and Iron shots are not available)
 water
IV. PROCEDURE

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
Safety is of paramount importance when conducting laboratory experiments, especially when
handling delicate equipment and potentially hazardous substances. To ensure the safety of the
experiment, we meticulously followed safety protocols. Prior to commencing the experiment, we
equipped ourselves with essential safety gear, including splash goggles, gloves, and laboratory
gowns. Additionally, our workspace was methodically organized to minimize the risk of
accidents. With these safety measures in place, we proceeded with the experiment.
We began our experiment by filling a 500 mL beaker with tap water, taking care to prevent
excessive splashing. We included a boiling chip in the beaker to facilitate a controlled boiling
process. Subsequently, the beaker was placed on a hotplate, and the temperature was set to
36°C. We patiently awaited the water's transition to the boiling phase.
Simultaneously, while some members of our team watched over the boiling beaker, others
prepared for the subsequent steps. This included pouring a precise quantity of stainless-steel
bearings into a test tube, filling it to the halfway point. As this task unfolded, we proceeded to the
area housing the balance. To ensure the accuracy of our measurements, we placed a weighing
boat, a Styrofoam cup in our case, onto the balance. We initiated the process by taring the
balance to 0.00g. Subsequently, the stainless-steel bearings were transferred to the weighing
boat, and their mass was determined. This recorded mass value was inscribed on line A of Table
4.1.
Upon concluding the mass measurement, the stainless-steel bearings were carefully transferred
into the test tube, which was destined for immersion in the boiling water within the beaker on the
hotplate. Prior to submerging the test tube, we attentively assessed the condition of the boiling
water in the beaker to confirm that it was at the appropriate temperature for the experiment.
Once satisfied, the test tube was inserted into the boiling water bath, using a test tube holder,
and allowed to equilibrate with the water for 10-15 minutes. Subsequently, we measured the
temperature of the boiling water bath with a precision of 0.1°C, and this value was duly recorded
on line B of Table 4.1.
Once the temperature of the stainless-steel bearings aligned with that of the boiling water bath,
we transitioned to the setup of the calorimeter. Utilizing a graduated cylinder, we measured
approximately 50.0 mL of cold tap water, with an accuracy of 0.1 mL, and added it to the
calorimeter. Upon complete transfer of the water, the calorimeter was sealed with its lid, and the
contents were gently swirled to ensure homogeneity. The temperature of the water within the
calorimeter was subsequently measured, and the recorded value, precise to 0.1°C, was
inscribed on line E of Table 4.1.
Post-collection of all essential data from the experiment, our workspace was diligently cleaned,
and the stainless-steel bearings were dried and stored for potential future use.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 Experiment 4
Specific Heat of Metal
REPORT SHEET
V. DATA AND OBSERVATION
Table 3.1: Heat of Fusion of Ice – observed and calculated data.
Item Stainless Steel Bearing
A. mass of shot, g 41.61 g
B. Temperature of water bath, ℃ 91.1℃
C. Volume of water, mL 91 mL
D. Temperature of water (initial), ℃ 21.1℃
E. Temperature of water (final), ℃ 27.2℃
F. Temperature Change (E-D), ℃ 6.1℃
G. Specific Heat , J/g ºC J
9.1503
g℃

Review Questions and Calculations

1. Using your experimental data for the lead shot and iron shot, calculate the specific
heat of lead and iron in J/g · °C and record your calculated values on line G of Table
4.1.
Q
C=
m× ∆ T
Q=mC ∆ T
mshot =41.61 g

mwater =91 mL=91 g

∆ T =6.1 ℃
J
Q=(91 g)(4.184 )(6.1 ℃)
g℃
Q=2322.5384 J
2322.5384 J
C=
(41.61 g)(6.1 ℃)
2322.5384 J
C=
253.821 g ℃
J
C=9.1503
g℃

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 2. The actual values for the specific heat of lead and iron are 0.127 J/g · °C and 0.450
J/g · °C, respectively. Calculate the percent error of the value you obtained
experimentally
9.1503−0.127
Percent Error=¿ ∨​×100 %
0.127
9.0233
Percent Error=¿ ∨​×100 %
0.127
Percent Error=|71.0496|× 100 %
Percent Error=7104.6063 %

3. Propose several possible explanations for experimental error.

1. Material Differences. One significant factor contributing to the experimental error

in my specific heat experiment was the use of stainless-steel bearings instead of
lead and iron shot. This material substitution introduced variations in the specific heat
because stainless steel has different properties than lead and iron.

2. Calorimeter Inefficiencies. I also observed that the calorimeter may not have
been perfectly insulated, leading to heat loss to the surroundings. The resulting
deviations in temperature measurements could have contributed to the error.

3. Incomplete Equilibration. It's possible that the hot stainless-steel bearings did
not completely equilibrate with the cold water in the calorimeter. This incomplete
equilibration could have resulted in an inaccurate representation of the specific heat
of the stainless steel.

4. Measurement Errors. Errors in measuring the mass of the stainless-steel

bearings, as well as inaccuracies in recording the initial and final temperatures and
the volume of water, may have affected the accuracy of my results.

5. Heat Loss. During the transfer of the hot stainless-steel bearings to the
calorimeter, I suspect that some heat might have been lost to the surroundings,
which could have influenced the overall heat transfer.

6. Contamination. I also considered the possibility of contamination, such as oil or

dirt on the stainless-steel bearings, affecting their heat capacity and potentially
causing errors in the specific heat determination.

7. Evaporation. There is a chance that water in the calorimeter may have

evaporated during the experiment, leading to an underestimation of the specific heat
of the stainless steel.

8. Uncertainty in Calorimeter Constants. I must account for the uncertainty in the

specific heat of the calorimeter (C_calorimeter), which could introduce some level of
error into the calculation.

9. Systematic Errors. I acknowledge that systematic errors in the equipment or

experimental procedure may also have played a role in the deviations observed in
my results. This might include inaccuracies in instrument calibration.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 10. Human Errors. It's important to note that human errors, such as inaccuracies in
reading instruments, not strictly following procedures, or making errors in data
recording, can collectively contribute to experimental inaccuracies.

In order to enhance the accuracy of my future experiments, I plan to address these

potential sources of error, conduct additional trials, and ensure that I use the correct
materials, such as lead and iron shot, as initially intended for a more precise
determination of specific heat.

4. After completing this laboratory, a student learned that his graduated cylinder
consistently delivered 5% less liquid than indicated. What effect did this error have
on the specific heat values that he determined experimentally? Why?
In light of the graduated cylinder consistently delivering 5% less liquid than indicated,
the specific heat values obtained in the experiment would be affected in the following
manner:

1. Underestimated Mass (m water ). It's important to recognize that the specific heat
(C) calculation depends on both the mass (m water ) and the temperature change (∆T)
of the water in the calorimeter. Since the graduated cylinder consistently delivers less
liquid, the mass of water used in the experiment is consistently underestimated. This
means that the actual amount of water in the calorimeter is higher than what was
measured.

2. Impact on Temperature Change (∆T). The underestimated mass also has an

effect on the calculation of the temperature change (∆T). As ∆T is determined by the
heat gained by the water, an underestimated mass results in the perceived
temperature change being greater than it actually is.

3. Overestimated Specific Heat (C). Given that specific heat (C) is inversely
proportional to both the mass (m water ) and the temperature change (∆T), the
underestimation of these values leads to an overestimation of the specific heat value.
This implies that the specific heat values determined experimentally would appear to
be higher than their actual values due to the systematic error in measuring the mass
of water.

In essence, the consistent error in the graduated cylinder, which causes it to deliver
less liquid than indicated, impacts the specific heat values by making them seem
higher than they really are. This is because the calculation is based on an
underestimated mass and temperature change due to the systematic error in the
measurement of the liquid volume.
VI. CONCLUSION AND RECOMMENDATION
Our experiment on determining the specific heat of stainless-steel bearings yielded valuable
insights and outcomes. Through a systematic and precise approach, we arrived at the
following conclusions:
1. Measurement Precision: The accuracy of our measurements played a pivotal role in
achieving reliable results. Small deviations in mass and temperature measurements
could introduce notable variations in the calculated specific heat. Therefore, it is

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1
 essential to employ calibrated instruments and meticulous measurement techniques to
reduce measurement-related discrepancies.
2. Calorimeter Efficiency: The well-insulated calorimeter proved to be a crucial
component in ensuring accurate results. The minimal heat exchange with the
surroundings, despite minor inherent losses, reinforced the experiment's reliability.
3. Heat Transfer Principles: Our experiment exemplified the fundamental principles of
heat transfer. The heat lost by the hot stainless-steel bearings equated to the heat
gained by the cooler water, allowing us to calculate the specific heat of the metal
accurately.
Based on our experiment, several recommendations emerge to enhance the accuracy and
precision of future investigations:
1. Measurement Protocol: To minimize measurement errors, employ instruments with the
highest precision and calibrate them regularly. Additionally, ensure measurements are
conducted meticulously, considering both the instruments' limitations and the conditions
under which the experiment is conducted.
2. Calorimeter Calibration: Periodic calibration of the calorimeter is essential to maintain
its efficiency. Routine checks and maintenance procedures should be executed to
uphold its insulation properties and reduce heat losses to the surroundings.
3. Data Analysis: Implement robust data analysis techniques, such as error propagation,
to account for uncertainties in measurements and calculations. A comprehensive
analysis framework can aid in better understanding and quantifying the associated
errors.
4. Multiple Trials: Conducting multiple trials of the experiment is advisable to mitigate
random errors and establish trends. A larger dataset enables the calculation of more
reliable average values and bolsters the validity of the conclusions.
In conclusion, our experiment to determine the specific heat of stainless-steel bearings
underscores the significance of precision, calorimeter efficiency, and heat transfer principles
in achieving accurate results. By adhering to the aforementioned recommendations, future
investigations can benefit from enhanced accuracy and rigor.

REFERENCES:
1. Serway, R. A., & Jewett, J. W. (2017). Physics for Scientists and Engineers with Modern
Physics. Cengage Learning.
2. Atkins, P., & de Paula, J. (2018). Atkins' Physical Chemistry. Oxford University Press.
3. Tipler, P. A., & Mosca, G. (2014). Physics for Scientists and Engineers. W. H. Freeman.
4. Helmenstine, A. M. (2020, November 2). This Is How to Calculate Percent Error.
ThoughtCo. https://www.thoughtco.com/how-to-calculate-percent-error-609584
5. Flowers, P. (2019, February 14). 1.5 Measurement Uncertainty, Accuracy, and Precision
- Chemistry 2E | OpenStax. https://openstax.org/books/chemistry-2e/pages/1-5-
measurement-uncertainty-accuracy-and-precision.

ALLEYA MAE B. FRANCISCO 10-26-2023

BSChE 1