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The Change in Temperature of Urea, Potassium Chloride, and Ammonium Nitrate When Reacting With Water

Amanda Conlon, Sara Nevedal, and Paige Redlin

Macomb Mathematics Science and Technology Center

Chemistry

Mrs. Hillard/ Mrs. Dewey/ Mr. Supal

May 21, 2015

Table of Contents

Introduction…………………………………………………………………...…………………...1

Review of Literature………..……………………………………………………………………..4

Problem Statement...……………………………………………………………………………..10

Experimental Design…………………………………………………………………………..…11

Data and Observations………………………………………………………………………..….13

Data Analysis and Interpretations………………………………………………………………..20

Conclusion……………………………………………………………………………………….28

Appendix A………………………………………………………………………………………34

Appendix B………………………………………………………………………………………35

Appendix C………………………………………………………………………………………38

Acknowledgements………………………………………………………………………………39

Work Cited……………………………………………………………………………………….40

Conlon-Nevedal-Redlin 1

Introduction

It’s the last quarter of the first game of the basketball season. The team’s all-star is about

to score the winning points on a breakaway layup when suddenly his momentum causes him to

trip resulting in a rolled ankle. As the trainer helps him off the court inspecting the injury, they

determine the ankle has suffered a sprain. This team could not possibly have a successful season

if their star player is unable to play for an extended amount of recovery time. In the U.S., about

30 million children and teens participate in some form of an organized sport and more than 3.5

million injuries each year are experienced by those participants. By far, the most common sports

induced injuries being sprains and strains. (“Sports Injury Statistics”)

Cold packs provide a quick and easy method of treatment for injuries like sprains to

reduce the longevity of the injury. Cold packs are practical due to their characteristics. Cold

packs can reduce bleeding into the tissues, swelling (inflammation), and muscle pain. These

qualities can lead to a shorter road back to full health and less likelihood of recurrence. (“Ice and

Heat Treatments for Injuries”)

Taking advantage of heat absorbing chemicals, the cold pack can drop in temperature

within a minute in a room temperature surrounding, unlike ice packs that take much longer due

to the fact that they need to be frozen hours beforehand.

The importance of this research can be discussed in many ways. The hope that the results

of this experiment could lead to establishments having the most effective cold pack on hand

when sudden injuries occur was the main motivation. Another example of the experiment’s

importance is shown in that instant cold packs can be used by paramedics because the immediate

cold temperatures are effective for emergencies. Due to the cold pack’s quick reaction time and

easily portable quality, paramedics find they are much more useful than conventional ice cubes


when it comes to tending to a patient in need of a cool down. Both athletic events and medical

emergencies could be benefit from further research in the endothermic reactions. If different

chemicals were tested to find endothermic reactions with different molarities, a scientific

breakthrough can be found. It is possible researchers could find ways to increase and advance the

technology currently used in cold packs.

Endothermic reactions are also found in more common applications than cold packs. A

process as simple as the melting of ice cubes are also considered an endothermic reaction. The

system absorbs the heat energy surrounding it, making the temperature of the surroundings or the

liquid the ice cubes were placed in drop. Therefore, endothermic reactions are the reason why ice

cubes are used to keep food and drinks cold.

The methods that were performed during experimentation included combining chemical

solvents and a water solute to create an endothermic reaction. The purpose of this process was to

determine which chemical produced the greatest temperature change; thus providing the most

effective choice in making cold packs. This would solve the issue of waiting for an ice pack to

freeze or having a cold pack that does not get cool quick enough due to a weak endothermic

reaction.

The chemicals ammonium nitrate, urea and potassium chloride were tested and the data

for temperature change collected was compared to discover the reaction creating the greatest

change in temperature. These chemicals were chosen because they are accessible to the general

public. Ammonium nitrate was chosen due to the fact that with previous research, it was found

this chemical is commonly found in cold packs. The two other chemicals were chosen to reveal

if there was a better solution to creating the cold packs. The same temperature of water was used,

with some fluctuation. This fluctuation represents environmental influences and the different


surroundings a cold pack can initially be kept in. This is due to the fact that if a cold pack was

made, in a hospital for example, the water would not have the same initial temperature each time.

The temperature change was found by recording the initial temperature of the water and

subtracting that from the lowest temperature reached during the reaction. From this process the

endothermic reaction with the greatest change in temperature would be revealed after analyzing

data therefore making the most effective cold pack.

By collecting data on the endothermic reactions between water with ammonium nitrate,

urea, and potassium chloride, the discovery of the best possible conditions needed to ensure the

most effective cold pack were found. This allows the general public and industries to be cost

effective and take full advantage of the chemicals and other tools that are purchased and used in

the multitude of industries and fields.


Review of Literature

No one wants to suffer from injuries like sprains, strains or bruises longer than necessary.

Instant cold or hot packs are often used by athletes and everyone alike to quickly reduce

inflammation and swelling. Early and effective treatment to injury can not only reduce pain but

take a faster path to healing by increasing blood flow. The quick fix to pain takes advantage of

chemicals that either absorb a lot of heat or release a lot of heat when dissolved in water (Hot

Pack/Cold Pack). When a chemical process absorbs heat it is called endothermic; when a

chemical process releases heat it is called exothermic. In cold packs, the chemical ammonium

nitrate, potassium chloride and urea are often used because they absorb a lot of heat when it

dissolves in water thus making the cold pack or its surroundings cold. In other words, these

chemicals dissolve in water endothermically. Water and the chemicals are kept in separate

compartments in the pack until the pack is needed. Then the chambers are broken and the

chemical dissolves in the water, absorbing heat and making the pack as cold as 0 degrees

Celsius. Cold packs generally last about 20 minutes. (“How Do Instant Hot and Cold Packs

Work?”)

Furthermore, it is known that chemical reactions proceed with the evolution or absorption

of heat. Through ice packs and heat packs, the absorption of heat can be felt. Enthalpy is the

amount of heat content used or released in a system at constant pressure. This heat flow

represents differences in chemical energy associated with the rearrangement of atoms in

molecules, the making and breaking of bonds to form new substances. This can also be

explained as a system losing its energy, while the surroundings are gaining energy. When

measured at constant pressure, this is the enthalpy change (ΔH) for the reaction. (“The


Thermodynamics of the Dissolution of Urea”) The ΔH of a reaction is negative if the process is

exothermic. The ΔH of a reaction is positive if the process is endothermic.

Figure 1. Enthalpy Diagrams

Figure 1 is an illustration of an exothermic and an endothermic reaction (“Enthalpy

Changes”). The first diagram is declining to show how energy or enthalpy is lost during an

exothermic reaction. The second diagram is increasing to show how energy is gained during an

endothermic reaction.

The chemicals used to create the endothermic reaction in cold packs can be referred to a

solute. A Solute is considered substance dissolved in another substance, usually the component

of a solution present in the lesser amount. For this experiment the three solutes would be

considered the ammonium nitrate, potassium chloride, and urea. A Solvent is considered the

component of a solution that is present in the greatest amount. It is the substance in which the

solute is dissolved. For this experiment the solvent would be considered the water each chemical

was placed in per trial. The dissolving of ammonium nitrate, potassium chloride, and urea into

water (H2O) would be classified as a dissolution. The term, dissolution, refers to a solute

dissolving in a solvent to form a solution. The solutions would be considered the substances left

in the calorimeter after the three types of reactions between the solvent and solutes have taken

place.


In order for the chemical reaction between the solvent and solute to measured accurately,

it must take place in a calorimeter. A calorimeter is an object used for calorimetry, or the process of

measuring the heat of chemical reactions. The calorimeter is part of the surroundings of the

experiment. The system is the part of the universe being studied, while the surroundings are the

rest of the universe that interacts with the system. There are also three types of systems, those

being open, closed and isolated. An open system is where both matter and heat can enter or leave

the system. A closed system is where matter cannot enter or leave the system, but heat can enter

or leave (“A System and Its Surroundings”). An isolated system is where neither matter nor heat

can enter or leave the system. A calorimeter would be an example of an isolated system, seen in

Figure 2 below.

Not allowing matter or heat to leave or enter results in receiving accurate measured data.

Another example often used to explain a system and its surroundings that a solvent and solute

can be commonly found in would be a thermos. The substance inside is the system, the thermos

itself is the system boundaries, and the atmosphere and world around the thermos is the

surroundings. For this experiment the surroundings is looked at because the purpose of the cold

pack is to cool the user down and the user would be a considered a surrounding.

Figure 2. Isolated System


Figure 2 shows an example of an isolated system (Pasquini). This container keeps the

energy and matter that could potentially be released from a reaction. The insulated system that

was used in this experiment was a calorimeter, which is pictured above. A calorimeter has an

inner and outer vessel which keeps it insulated.

When the chemical reaction between the solute and solvent happens inside a calorimeter,

the process of equilibrium takes place. The equilibrium constant (K) defines the relationship

among the concentrations of chemical substances involved in a reaction at equilibrium. The Le

Châtelier's principle states that if a stress, such as changing temperature, pressure, or

concentration, is inflicted on an equilibrium reaction, the reaction will shift to restore the

equilibrium. For exothermic and endothermic reactions, this added stress is a change in

temperature. The equilibrium constant shows how far the reaction will progress at a specific

temperature by determining the ratio of products to reactions using equilibrium concentrations.

Considering the fact that the only type of stress that changes the equilibrium constant (K) is

temperature is not changed in the experiment, the equilibrium constant (K) will not change

throughout the trials.

To understand how the reaction takes place it is needed to look at the two driving forces

of the chemical processes, energetics and entropy. These factors determine whether a change

occurs in the system and how energy flows if it does. The chemistry in energy deals with

attractive and repulsive forces between particles at the molecular level. All of the trillions of

molecules are constantly moving, vibrating, and rotating at different rates. It can be thought that

temperature is a measurement of the average motion or Kinetic Energy of all of these particles;

with an increase in movement creating an increase in temperature and vice versa.


The flow of heat in any chemical transformation depends on the relative strength of

particle interactions in each of the substances chemical states. When particles have a strong

mutual attraction force they move rapidly towards one another until they get so close that

repulsive forces push them away. If the additional attraction was strong enough, particles will

keep vibrating back and forth in this way. The stronger the attraction will create a faster

movement; since heat is essentially motion, substance changes to a state in which these

interactions are stronger. Cold packs do the opposite; which means that when a solid dissolves in

water the new interactions of solid molecules and water molecules with each other are weaker

than separate reactions that existed before. This makes both types of particles slow down, on

average cooling the whole solution. The substance changed to a state where the interactions are

weaker due to entropy.

Entropy describes how objects and energy are distributed based on random motion. An

example of entropy would be the air in a room and all of its unique possible arrangements of

molecules. There is a possibility that all of the oxygen could be grouped together and all of the

nitrogen were to be grouped in another area. It is often found that air will have nitrogen and

oxygen mixed together. This is why air is typically presented in this state. If there is strong

attractive forces between particles, the probability of some configuration can change even to the

point where the odds don't favor certain substances mixing. An example would be oil and water.

In the case of ammonium nitrate, urea, or potassium chloride used in cold packs,

attractive forces are not strong enough to change the odds of grouping. Therefore, random

motion made the particles composing of the solid separated by dissolving into the water and

never returning to its solid-state. To put in simply, the cold packets get cold because random

motion creates more configurations where the solid and water are mixed together. All of these


interactions have even weaker particle interaction. Thus having less overall particle movement

and less heat then there was inside the unused pack.


Problem Statement

Problem:

Will potassium chloride KCl, urea CO(NH2)2, or ammonium nitrate NH4NO3 produce the most

effective instant cold pack with the greatest change in temperature when mixed with water H2O?

Hypothesis:

If potassium chloride KCl, urea CO(NH2)2, and ammonium nitrate NH4NO3 are mixed with water,

the ammonium nitrate NH4NO3 reaction will have will have the greatest change in temperature

when mixed with water, making it the most effective chemical for an instant cold pack.

Data Measured:

Various chemicals were mixed with water in a calorimeter. The independent variable is

the reactant used to be mixed with water. The reactants are potassium chloride, urea, and

ammonium nitrate, each trial contains 0.25 moles of one chemical. The dependent variable is the

change in temperature measured in Celsius. The change in temperature is found by subtracting

the lowest temperature from the initial temperature.


Experimental DesignMaterials:Calorimeter 325 g of Ammonium Nitrate, NH4NO3

TI-nspire calculator 250 g of Urea, CH4N2O150 mL Graduated Cylinder 300 g of Potassium Chloride, KClLabQuest Scale (To the nearest 0.0001g)(90) Weigh Boats LabQuest Thermometer Probe (0.1 C)

Procedures: 1. To avoid bias, before starting experimentation randomize the order of the trials using the

randomize function on a TI-nspire calculator (Appendix A).

2. Measure out the correct amount for the chemical for the trial, using 7.5 grams for urea, 9.25

grams for potassium chloride and 10 grams for ammonium nitrate.

3. Measure 150 mL of water using the graduated cylinder.

4. Set up the LabQuest with the thermometer apparatus adjusted to measure temperature as a

function of time for 180 seconds.

5. Pour the 150 mL of water in to the calorimeter.

6. Insert the thermometer probe into the calorimeter and press the play button to begin recording

the temperature.

7. After 5 seconds, pour the chemical being used into the calorimeter and begin stirring with

temperature probe.

8. Observe the graph of time and temperature and after 180 seconds, remove the thermometer

probe.

9. Record the initial and lowest temperature reached in Celsius, observed from the LabQuest.

10. Safely dispose of the contents in the calorimeter, then wash it thoroughly and dry it with

paper towel.

12. Repeat steps 1-11 thirty times for each chemical.


Figure 3. Setup of LabQuest and Calorimeter

Figure 3 demonstrates how set up a trial for the experiment. The thermometer probe is

connected to the LabQuest then put into the calorimeter, where the reactants are. The

thermometer and the LabQuest will record the change in temperature and the time. Properly

setting up the LabQuest will allow accurate data (Appendix C).


Data and Observations

Data:

Table 1 Urea Reaction Temperature Changes

Trial

Initial Temperature

of Water(°C)

Lowest Temperature

Reached(°C)

Change in Temperature

(°C)

1 20.5 17.6 -2.92 20.6 17.4 -3.23 20.3 17.2 -3.14 20.1 17.2 -2.95 20.8 17.9 -2.96 22.9 20.4 -2.57 21.7 18.8 -2.98 22.0 19.1 -2.99 21.1 18.1 -3.010 20.6 17.7 -2.911 23.6 20.6 -3.012 21.4 19.3 -2.113 21.1 18.0 -3.114 23.5 20.2 -2.315 22.7 19.1 -3.616 22.0 18.8 -4.217 21.5 18.4 -3.118 21.2 17.6 -3.619 20.4 17.9 -2.520 20.9 18.5 -2.421 20.9 18.4 -2.522 21.1 18.5 -2.623 20.0 17.8 -2.224 20.1 18.5 -1.625 22.0 20.0 -2.026 21.8 18.8 -3.027 21.5 18.2 -3.328 21.1 17.5 -3.629 20.9 17.6 -3.330 20.8 18.3 -2.5

Average 21.3 18.2 -3.1


Table 1 shows the data collected when 7.5 grams of urea were mixed with water. This

table shows the initial temperature recorded for each trial, along with the lowest temperature

reached during the reaction. The table also shows the change in temperature for each trial which

is found by subtracting the initial temperature from the lowest temperature. The averages for the

initial temperature, the lowest temperature, and the change in temperature are included in the last

row.

Table 2Potassium Chloride Reaction Temperature Change

Trial

Initial Temperature

of Water(°C)

Lowest Temperature

Reached(°C)


(°C)

1 19.8 16.1 -3.72 20.4 16.7 -3.73 20.2 16.6 -3.64 20.9 17.3 -3.65 21.0 17.3 -3.76 20.9 17.2 -3.77 21.1 17.5 -3.68 21.1 17.4 -3.79 19.9 16.4 -3.510 19.8 16.3 -3.511 19.8 16.2 -3.612 20.2 16.3 -3.913 19.9 16.3 -3.614 19.5 16.2 -3.315 20.9 17.5 -3.416 20.6 16.8 -3.817 19.8 16.5 -3.318 19.3 16.1 -3.219 19.8 16.3 -3.520 19.7 16.4 -3.321 20.0 16.8 -3.222 19.9 16.5 -3.423 19.4 16.3 -3.224 19.8 16.5 -3.325 20.2 16.7 -3.5


Trial

Initial Temperature

of Water(°C)

Lowest Temperature

Reached(°C)


(°C)

26 20.1 16.8 -3.327 20.2 16.7 -3.528 19.7 16.5 -3.229 19.5 16.2 -3.330 20.7 16.9 -3.8

Average 20.1 16.6 -3.5Table 2 shows the data collected when 9.25 grams potassium chloride was used. This

table shows the initial temperature recorded for each trial, along with the lowest temperature

reached during the reaction. The table also shows the change in temperature for each trial which

is found by subtracting the initial temperature from the lowest temperature. The averages for the

initial temperature, lowest temperature, and the change in temperature are included in the last

row of the table.

Table 3 Ammonium Nitrate Reaction Temperature Change

Trial

Initial Temperature of

Water(°C)

Lowest Temperature

Reached(°C)


(°C)

1 20.9 15.4 -5.52 20.1 16.6 -3.53 21.6 17.8 -3.84 22.3 17.4 -4.95 22.1 16.9 -5.26 21.9 18.5 -3.47 22.6 17.9 -4.78 22.1 17.5 -4.69 21.7 17.4 -4.310 22.1 17.4 -4.711 24.2 19.8 -4.412 25.7 21.3 -4.413 24.8 20.5 -4.314 23.3 19.6 -3.715 22.6 18.1 -4.5

Trial Initial Lowest Change in


Temperature of Water(°C)

Temperature Reached

(°C)

Temperature (°C)

16 22.0 18.8 -4.217 21.5 17.4 -4.118 21.2 16.6 -4.619 21.4 17.9 -3.520 20.9 17.3 -3.621 21.9 17.9 -4.022 21.1 16.5 -4.623 20.0 16.8 -3.224 20.1 16.4 -3.525 22.0 18.1 -3.926 21.5 16.9 -4.627 20.6 16.5 -4.128 23.2 19.3 -3.929 22.1 18.2 -3.930 21.8 17.6 -4.2

Average 22.0 17.8 -4.2Table 3 shows the data collected when 10 grams ammonium nitrate was used. This table

shows the initial temperature recorded for each trial, along with the lowest temperature reached

during the reaction. The table also shows the change in temperature for each trial which is found

by subtracting the initial temperature from the lowest temperature. The averages for the initial

temperature, the lowest temperature, and the change in temperature are included in the last row

of the table.

Table 4Urea Reaction ObservationsTrial Observations

1 Researcher C measured, Researcher B stirred. Water turned cloudy immediately, then at about 100 seconds turned clear again.

2 Researcher C measured, Researcher B stirred. Water turned cloudy immediately, took a little over 100 seconds to turn clear again, got a little water around top of cal.

3 Researcher C measured, Researcher B stirred. Still small particles after 100 seconds.4 Researcher C measured, Researcher B stirred. Cloudiest water yet.5 Researcher C measured, Researcher B stirred. Cloudy to clear at 120 seconds.6 Researcher C measured, Researcher B stirred. Some air bubbles while stirring

Trial Observations7 Researcher C measured, Researcher B stirred. Consistently clear, temperature went up 1


degree at first, then dropped about 30 seconds in.8 Researcher C measured, Researcher B stirred. Water cleared earlier than 100 seconds.

9 Researcher C measured, Researcher B stirred. Spilled a little water around rim, cloudy for longer than 100 seconds.

10 Researcher C measured, Researcher B stirred. Turned clear quicker than 100 seconds.

11-12 Researcher C measured, Researcher B stirred. Urea was more powdery since it was bottom of tub, never very cloudy, cleared up faster than 100 seconds.

13 Researcher C measured, Researcher B stirred. Urea was more powdery since it was bottom of tub, very cloudy, cleared up after 100 seconds

14Researcher C measured, Researcher A stirred. Less cloudy, went clear at 100 seconds. Urea from first container ran out, the new chemical looks slightly different and was used for the rest of the trials. Turned clear at 100 seconds.

15 Researcher C measured, Researcher A stirred. Urea was clumpy before starting reaction. Turned clear at 100 seconds.

16- 17

Researcher C measured, Researcher A stirred. Turned clear before 100 seconds.

18 Researcher C measured, Researcher A stirred. Still slightly cloudy after 100 seconds.19-21 Researcher C measured, Researcher A stirred. Solution was overall very cloudy.

22 Researcher C measured, Researcher A stirred. As stirring occurred, a small amount of water splashed onto the top rim of the vessel. Turned clear, shortly after 100 seconds.

23 Researcher B measured, Researcher A stirred. Solution turned clear after 100 seconds.

24 Researcher B measured, Researcher A stirred. Solution started turning clear before 100 seconds.

25 Researcher B measured, Researcher B stirred. Turned clear before 100 seconds.

26 Researcher B measured, Researcher B stirred. Slightly clear around 60 seconds, completely clear around 100 seconds.

27 Researcher B measured, Researcher A stirred. Solution started turning clear before 100 seconds.

28 Researcher B measured, Researcher A stirred. Initial urea appears to be clumpy, less fine crystals than usual. Turned clear before 100 seconds.

29Researcher A measured, Researcher C stirred. A few clumps were found in the initial urea. Turned clear before 100 seconds. The clump in the chemical took longer to dissolve.

30 Researcher A measured, Researcher C stirred. Very CloudyTable 4 contains the observations that were made for each trial of the urea and water

reactions. As trials were conducted, observations were made and recorded in this table.

Table 5Potassium Chloride Reaction Observations


Trial Observations1 Researcher A weighed, Researcher C stirred. Only slightly cloudy.

2 Researcher A weighed, Researcher C stirred. Slightly cloudy, some bubbles around sides.

3-7 Researcher A weighed, Researcher C stirred. Clear at 70 seconds.8 Researcher B weighed Researcher C stirred, large chunks in powder.9 Researcher B weighed, Researcher C stirred, clear at 70 seconds.10 Researcher B weighed, Researcher C stirred, water splashed.

11 Researcher A weighed, Researcher C stirred, clear at 70 seconds, some chloride got onto top part of calorimeter.

12 Researcher A weighed, Researcher C stirred, some chunks in the powder.13-14 Researcher C weighed, Researcher A stirred, clear at 70 seconds.

15 Researcher A weighed, Researcher B stirred. Clear at 75 seconds.16 Researcher A weighed, Researcher B stirred. Clear at 80 seconds.17 Researcher A weighed, Researcher B stirred. Clear at 70 seconds.18 Researcher A weighed, Researcher B stirred. Clear at 75, water splashed out.19 Researcher A weighed, Researcher B stirred. Clear at 75.20 Researcher A weighed, Researcher B stirred. Clear at 80, some particles left at bottom.21 Researcher A weighed, Researcher B stirred. Clear at 85 seconds.22 Researcher C weighed, Researcher B stirred, clear at 75 seconds.23 Researcher C weighed, Researcher B stirred. Clear at 80 seconds.

24-27 Researcher C weighed, Researcher B stirred, clear at 80 seconds.28 Researcher C weighed, Researcher A stirred, clear at 70 but some particles remained.29 Researcher C weighed, Researcher A stirred, clear at 75 seconds.

30 Researcher C weighed, Researcher A stirred, clear at 75 seconds. There was a big clump in the chemical before it was put into the reaction.

Table 5 contains the observations that were made for each trial of the potassium chloride and water reactions. As trials were conducted, observations were made and recorded in this table.

Table 6Ammonium Nitrate Reaction ObservationsTrial Observations

1-3 Researcher C stirred, Researcher A weighed. Very cloudy. Cleared up around 150 seconds.

4 Researcher C stirred, Researcher A weighed. Some small particles remained after 100 seconds.

5-7 Researcher C stirred, Researcher A weighed. Very cloudy. Cleared up around 150 seconds.

8 Researcher C stirred, Researcher A weighed. Very cloudy. Cleared up around 115 seconds.

9 Researcher B stirred, Researcher A weighed. Clear at 120 seconds.Trial Observations10-15 Researcher B stirred, Researcher A weighed. Cloudy, clear at 110 seconds.


16-30 Researcher B stirred, Researcher A weighed. Cloudy, clear at 100 seconds.Table 6 contains the observations that were made for each trial of the ammonium

nitrate reactions. As trials were conducted, observations were made and recorded in this table.


Data Analysis and Interpretation

The data collected for this experiment involves the temperature change taking place

during a chemical reaction. This reaction is created by dissolving the chemicals urea, ammonium

nitrate and potassium chloride in water inside a calorimeter. Throughout the duration of our

experimentation steps were taken to ensure that the data was reliable and valid. One step taken to

insure validity was to eliminate bias in the results by randomizing trials. This experiment

involved 3 chemicals with 30 trials each, resulting in 90 trials total. Randomization was

accomplished this by using a Ti-Nspire CX calculator, explained in Appendix A. Another step

involved replication as a method used to create valid data. This was done by seeing the same

amount of water, correct amount of chemical substance and the same calorimeter was used for

each trial. Furthermore, all of the trials were conducted with the same procedure and setting.

Also because our experiment involved temperature change, the correct amount of time was taken

for and in between each trial, ensuring previous trials would not have effect on future trials.

Based on the design of experiment and the collected data, it was decided the most

appropriate analysis was descriptive. A T-test or Z-test was not found an appropriate test

because there is no known value to compare the data to. A DOE was not appropriate because

only one factor being temperature was being tested, where a DOE needs a variability of at least 2

or 3 factors to be applicable. Due to this experiment’s criteria not meeting the requirements for

these analytical tests, it was decided descriptive would be the most suitable. Descriptive analysis

was carried out by using probability plots to show the normality of the data and box plots to

show spread of the data.


Figure 4. Normal Probability Plot of Temperature Change During Urea Reactions

The data shown in the normal probability plot for the temperature change during urea

reactions seems reasonably normal. All plotted values fall close to the line and are within a small

range of 1.1°C from -3.6°C to -2.5°C. There are no outliers and the closely stacked dots that

seem to grow farther from the line simply show repeating values.

Figure 5. Normal Probability Plot of Temperature Change During Potassium Chloride Reaction

The data shown in the normal probability plot for the temperature change during

potassium chloride reactions seems reasonably normal. All plotted values fall close to the line

and are within a small range of 0.7°C from -3.9°C to -3.2°C. There are no outliers and the


closely stacked dots that seem to grow farther from the line simply show repeating values.

Figure 6. Normal Probability Plot of Temperature Change During Ammonium Nitrate Reactions

The data shown in the normal probability plot for the temperature change during

ammonium nitrate reactions seems reasonably normal and contains no outliers. All plotted values

fall close to the line and are within a small range of 2.3°C from -5.5°C to -3.2°C.

Figure 7. Box Plot of Temperature Change for Urea Reactions

Figure 7 shows a boxplot of the temperature change calculated per urea trial. The plot

shows that the data fell between -2.5°C and -3.6°C, giving the data set a range of 1.1°C. The data

also appears to be not skewed at all and there are no outliers which suggest that the trials were


very consistent. In addition to the lack of skew and outliers, the plot also shows that the median

experimental temperature change, -3.1°C, is very close to the exact same of the mean value, -

3.1°C, the average of all 30 trials.

Figure 8. Box Plot of Temperature Change for Potassium Chloride Reactions

Figure 8 shows a boxplot of the temperature change calculated per potassium chloride

trial. The plot shows that the data fell between -3.2°C and -3.9°C, giving the data set a range of

0.7°C. The data also appears to not be skewed, hence there are no outliers which suggest that the

trials were consistent. The plot also shows that the median experimental temperature change, -

3.5°C, is the exact same of the mean value, -3.5°C. This shows that the data is very normal

because the median is a resistant value which median that if there was abnormal data then it

would be not affected. On the other hand the mean is a nonresistant value which means that if

there is an abnormal data point then it would be affected. The fact that the mean and the median

are the same due to one being resistant and one being nonresistant shows that the data is very

normal.


Figure 9. Box Plot of Temperature Change for Ammonium Nitrate Reactions

Figure 9 shows a boxplot of the temperature change calculated per ammonium nitrate

reaction trial. The plot shows that the data fell between -3.2°C and -5.5°C, giving the data set a

range of 2.3°C. This has the highest range out of the three reactions suggesting that the chemical

of ammonium nitrate may have not produced reactions with consistent results, as it is felt human

error during trials was avoided. In addition to the lack of outliers, the plot also shows that the

median experimental temperature change, -4.2°C, is the exact same as the mean value, -4.2°C,

the average of all 30 trials. This shows that the data is very normal because the median is a

resistant value which means that if there was abnormal data then it would not be affected.

Conversely, the mean is a nonresistant value which means that if there is an abnormal data point

then it would be affected. The fact that the mean and the median are the same due to one being

resistant and one being nonresistant shows that the data is very normal.

During the course of the trials, specific heat was used to compare the trials’ validity. The

known values were plugged into the SM ΔT=SM ΔT equation and then solved for S (Appendix

B). These values were then taken and plugged into a formula to find the correction factor.

Even though the correction factor seems necessary for this analysis, it is not significant to

the data and what the experiment was conducted to find. The experimental specific heats and


actual specific heats were compared, simply to justify the data. The main reason this is not

significant is because this experiment was conducted to measure the total change in temperature,

not the highest or lowest specific heat. This test is being done simply to justify the normality of

the data and compare it to the established specific heats for these chemicals.

For each chemical, the average high and low temperature along with the average

temperature of the room were taken and plugged into the SM ΔT=SM ΔT formula. (Appendix

B). After the calculations, the average experimental specific heat ranged from potassium

chloride’s value of 3.9*101 J/g°C to urea’s value of 4.4*101 J/g°C. When compared to the

researched specific heats, not only were the numbers off by a significant amount, but the order

was also significantly different. For example, the order from highest specific heat to lowest

specific heat in the experimental specific heats went in the order from urea, ammonium nitrate,

to potassium chloride. This is quite different from the order of the researched specific heats

which went in order from the highest being potassium chloride, to urea, to the lowest researched

specific heat being the ammonium nitrate value of 4.6*103 J/g°C.

After calculating the specific heat for each of the three chemicals, it was determined that

the data did not seem accurate. It was then decided to calculate a correction factor and apply it to

the experimental specific heats to calculate a new and more accurate specific heat to compare to

the researched specific heat. By using the correction factor formula seen below for each of the

three chemicals, a correction factor was found. These correction factors range from the lowest

percentage of 86.92% for the urea trials to the highest correction factor of 91.02% for the

potassium chloride trials. All of these correction factors are significantly high, but the trials that

were conducted for potassium chloride had the highest correction factor. This means that if the

experimental specific heat is multiplied by the correction factor, a new specific heat is found that


is more accurate. These three values ended up being 3.8*103 J/g°C for the urea trials, 3.6*103

J/g°C for the potassium chloride trials, and 3.8*103 J/g°C for the ammonium nitrate trials. Even

with the correction factor, these three values are not exactly the same as the researched specific

heat for each chemical.

Table 7Correction Factors of Urea, Potassium Chloride, and Ammonium Nitrate

Initial Specific Heat

(J/g°C)

Correction Factor

(%)

Specific Heat with Correction Factor

(J/g°C)

Actual Specific Heat

(J/g°C)

Urea 4.4*101 86.92 3.8*103 5.6*103

Potassium Chloride

3.9*101 91.02 3.6*103 5.9*103

Ammonium Nitrate

4.3*101 88.11 3.8*103 4.6*103

Table 7 shows the correction factors found in previous figures put to use. The average

specific heat found from each chemical was placed in the first column. The next column was

filled with the correction factor for each chemical. The next column labeled ‘Number with

Correction Factor’ was found by multiplying the initial specific heat by the correction factor as

an integer, not a percent. It was not multiplied by a percent or .8692 for example, due to the fact

that by multiplying the experimental specific heat by a decimal a smaller number than the initial

specific heat would be calculated. The last row simply compares the new number with the

correction factor implemented to the actual specific heat of these chemicals. Notice how with the

correction factor, there is a significant difference in the comparison of the experimental specific

heat to the actual specific heat.

While conducting our experiment, it was believed that the correction factor would end up

low, considering that all of the data seemed to be normal at the time. A low correction factor


would have justified our data as reliable. Even though our correction factors are high, usually

indicating low normality in the data, other evidence like the normal probability and box plots

justifies the data.


Conclusion

When the chemicals urea, potassium chloride, or ammonium nitrate are mixed with water

they reduce the temperature of the water, thus the reason why these chemicals are used in cold

packs. Cold packs are often used in athletics for treatment of injuries like sprains and strains.

Cold packs make reducing symptoms of these injuries such as inflammation and swelling more

efficient. The overall purpose of this experiment was to determine which chemical produced the

greatest temperature change and therefore provides the most effective choice in making a cold

pack.

The original hypothesis in this experiment was that if potassium chloride, KCl, urea,

CO(NH2)2, and ammonium nitrate, NH4NO3, are mixed with water, the ammonium nitrate NH4NO3

reaction will have the greatest change in temperature. Hence making it the most effective

chemical for an instant ice pack. The results showed NH4NO3 as having the greatest change in

temperature, with an average change of -4.2°C. This data results in the hypothesis failing to be

rejected. These research results support the current work in the field of ice packs. Ammonium

nitrate is the most widely substance used to react with H2O because unlike other endothermic

reactions, the mixing of ammonium nitrate and water has proven to absorb significantly more

energy than other chemicals.

These results occurred due to how an endothermic reaction is carried out. The reaction is

absorbing heat, so the heat in the solution is being consumed by the reaction. Losing heat

therefore made the surroundings of the chemical solute and H2O solvent solution become colder.

In order for a chemical reaction to take place, the reactants collide. The reaction took the

available heat and changed it into chemical bonds. The collision between the molecules in the

H2O solute and chemical solvent reaction provides the kinetic energy needed to break the


necessary bonds so that new bonds can be formed. This process uses heat energy, therefore

making the surroundings of the reaction’s temperature drop, due to the absence of heat.

In the case of ammonium nitrate, urea, or potassium chloride used in cold packs,

attractive forces are not strong enough to change the odds of the same types of particles being

grouped together. Therefore, random motion made the particles composing of the solid separated

by dissolving into the water and never returning to its solid state. This has to do with collision

theory. The collision theory is based on the assumption that for a reaction to occur it is necessary

for molecules to come together or collide with one another. Not all collisions, however, bring

about chemical change. The molecules must be oriented in a manner favorable to the necessary

rearrangement of atoms and electrons. (“Collision Theory”) Put simply, the cold packs get cold

because random motion creates more configurations where the solid and water are mixed

together. All of these interactions have even weaker particle interaction. Thus having less overall

particle movement and less heat then there was inside the unused pack.

Due to our controls and secure experimental set up, the data was accurate and reliable.

This data showed that he reaction between ammonium nitrate and water took the largest amount

of energy to occur. Therefore, this indicates that the system gained the greatest amount of

energy, having the greatest temperature change. This made the surroundings of the ice packs, the

part that can be felt by the consumer, drop in temperature becoming cold to the touch. The

surroundings give the greatest amount of energy up to the system in the ammonium nitrate

reaction compared to the potassium chloride and urea reaction. Ammonium nitrate is also a

practical chemical to use for instant cold packs considering the price. A homemade cold pack can

be created with ammonium nitrate for under $5.00, this supports the fact that it is easy for every

sporting event to have these instant cold packs on hand in case of emergencies.


The data and conclusion of this experiment match those of previous research. After

testing endothermic reactions of ammonium nitrate, ammonium chloride, and potassium

chloride, Cadrette-Jaworski determined through their data that ammonium nitrate created the

highest temperature change. They also found that it was close to the fastest reaction rate, coming

to the conclusion that ammonium nitrate is the choice chemical for efficient cold packs

(Cadrette). The fastest reaction rate in Chadrette-Jaworski’s research was regarding the least

amount of time for the reaction to reach its’ lowest temperature. Through the data collected in

this experiment, it was also found that ammonium nitrate has the greatest temperature change.

Comparing the results from this experiment to previous research, it can be concluded that this

research and Cadrette-Jaworski’s research are comparable in results.

Throughout the experiment, the experimental design that was followed was determined to

be well constructed. With the use of randomization (Appendix A), constants, and careful

precision while performing trials, the experiment itself was considered to be constructed well by

the scientists behind it. By conducting the randomization of trials, no two of the same trials were

consecutively conducted. This prevents the data from becoming routine. Constants were found in

the experiment to increase validity. Keeping the same calorimeter, water faucet, temperature

probe, and measuring tools can be considered constants that prove that a well-structured and

reliable procedure was followed. The last element that lead to having a well-constructed

experiment would be careful precision while performing trials. While measuring chemical

solvents and water solute, careful precautions were taken to measure as close as to the exact

weight or volume as possible. These elements resulted in this experiment being reliable, thus

creating accurate results.


Even though all procedures were done to maintain the validity of the experiment there

were still minor errors when the data was analyzed statistically. For each chemical the

experimental specific heat was around ten times smaller than the actual researched specific heat.

This is determined to be an undetected error by the researchers. To account for inconsistency of

the data a correction factor was calculated for each chemical using a formula (Appendix B). The

high correction factor is determined as insignificant after considering a few factors. The Central

Limit Theorem is reached (Appendix B), the boxplots are analyzed as normal, and all of the

trials’ specific heats were equally different from the actual specific heats. These three factors are

indications that the data concerning specific heat and the high correction factor was proved to be

insignificant.

The overall results in this experiment recorded for the changes in temperature for every

chemical had a relatively small spread with a range from -2.5°C to -5.5°C. Individually,

ammonium nitrate had the largest range of temperature change of 2.3°C from -3.2°C to -5.5°C.

This range, being the largest, indicates the possibility that there are large differences between

individual scores. Large ranges can be due to a lurking variable. One possible lurking variable

was the temperature of the surroundings, which was unaccounted in the experiment and unable

to be constantly controlled. The temperature of the water used in the calorimeter was not

constant, which affected the reaction rate because temperature change affects the kinetic energy

of molecules and different reactants’ rate of contact with each other. Also, when the solvent was

added to the solute, five seconds were allowed to pass before starting to mix the solution

together. The exact time of the five second waiting periods could fluctuate by a small amount,

therefore the rate had a chance of being inconsistent from trial to trial. Yet another source of


error is that often when a solute was poured into the solvent to make the solution, some solute

affixed to the wall of the container and had a delayed reaction with the solvent.

These errors can be fixed in many possible ways. Errors could be prevented if room

temperature distilled water was used in place of tap water with fluctuating temperature. The

water could then be considered a constant, not having to worry about the different reaction rates

in different temperatures of water. The error of inconsistency of the reaction rate could be

avoided by using a timer to ensure the waiting period is the same for every trial. Another way to

reduce errors with reaction rate could be to find a more efficient device to pour the chemical into

the calorimeter. The weigh boat is made of a flexible plastic that must be physically bent to pour

and sometimes resulted in the chemical substance not going directly into the water. It could be

possible the use of a pipe or tube to transport the chemical from the weigh boat to the calorimeter

would result in all of the substance going directly in the water and avoiding sticking to the sides.

In addition to correcting these errors, there are other potentials for further research in this

field. There are many possible factors that researchers could explore. Some factors that could be

tested are a greater array of chemicals, initial water temperature, and changing the previous

chemicals used to common chemicals found in heat packs. Testing more chemicals, such as the

reaction of baking soda and citric acid, could possibly uncover a chemical that has a greater

temperature change that has not yet been discovered to possess this quality (“Home Chemistry”).

Thus, this would lead to a more efficient instant cold pack. Also, initial water temperature could

be researched more thoroughly to see what effect it has on the reaction. A high or low water

temperature could have an effect on the range of the temperature change. The same procedure

used for this experiment could easily be applied to test heat packs by changing the chemicals that

were tested. It could be researched to see if exothermic reactions have around the same amount


of temperature change as the endothermic reaction that is found in cold packs by comparing the

two experiments absolute value.

A possible application relating to endothermic reactions would be instant cold packs.

Many injuries go untreated during sports due to ice not being readily available. If one does not

apply a cold compress after the initial trauma, an injury can easily become worse due to swelling

and inflammation. A cold pack is a portable, easily accessible method of treatment to prevent

these symptoms. By taking advantage of endothermic reactions, an instant cold pack contains a

chemical that absorbs heat when mixed with water. An effective instant cold pack should be able

to have a significantly large temperature change. A certain chemical reaction that shows a

significant change or drop in temperature could be used to make an improved cold pack. Finding

a more efficient and affordable way to start injuries on the path to quick recovery could lead to

less pain for athletes in their futures. Injuries that are not treated properly are more likely to re-

occur, so the hope would be if there were an easy way to achieve treatment the use of it would

increase and time out on injury would decrease.


Appendix A: Randomization

Materials:

Ti-Nspire CX calculator

Procedure:

1. Turn on calculator and open a calculator page.

1. Press the menu button.

2. Scroll down and choose the “5: Probability” by pressing enter.

3. Scroll down and choose the “4: Random” option by pressing enter.

4. Select the “2: Integer” function and press enter.

5. The display of “randInt ()” should be shown on the screen. Enter the minimum number

for the range of numbers desired for randomization. For example, if the numbers 1

through 50 are to be randomized, enter 1.

6. Enter a comma and then the last number desired for the range of numbers that the

randomization includes. For example, in the randomization of 1 through 50, 50 would

now be entered.

7. Press enter. A number will appear, representing a trial number.

8. Continue pressing enter until all trials are accounted for. If a number repeats, do not take

it into account.

9. Repeat steps 1-9 for each run.


Appendix B: Descriptive Data and Analysis Equations

Throughout the course of experimentation, two main formulas were used when analyzing

the data. These two formulas were used to find the specific heat while the second formula was

used to find a correction factor. Data from the experiment that was conducted was taken and

plugged into these formulas so desired values could be found.

SM ΔT=SM ΔT

Shown in figure 1 below is sample work of the specific heat formula.

Urea Trials Sample Equation

SM ΔT=SM ΔT

4.184(150.0)(18.2-20.7) = S(7.5)(18.2-23.0)

4.184 (150)(18.2−20.7)(7.5)(18.2−23.0)

=S

4.4*101 kj/mol=S

Potassium Chloride Trials Sample Equation

SM ΔT=SM ΔT

4.184(150.0)(16.6-20.3)= S(9.25)(16.6-23.0)

4.184 (150)(16.6−20.3)(9.25)(16.6−23.0)

=S

3.9*101 kj/mol=S

Ammonium Nitrate Trials Sample Equation

SM ΔT=SM ΔT

4.184(150.0)(17.8-21.4)= S(10.0)(17.-23.0)

4.184 (150)(17.8−21.4)(10.0)(17.8−23.0)

=S


4.3*101 kj/mol=S

Figure 1. Specific Heat Formula.

In the formula above each variable has a different value plugged in when solving. On the

left side of the formula, S stands for the specific heat of water. The specific heat of water is used

because our reactions were performed by dissolving the chemicals in water. On the left side of

the formula, M stands for the mass of water in grams we used which was the same in every trial.

Lastly on the left side of the equation, ΔT stands for change in temperature of the water as the

reaction with the chemical took place. The change in temperature was found by subtracting the

lowest temperature reached during the reaction from the initial temperature of the water. As for

the right side of the equation, the S stands for specific heat of the chemical used, which is what

we are using the above formula to solve for. On the right side of the formula, the M stands for

the mass of the chemical in grams used in the reaction. Lastly on the right side of the formula ΔT

stands for change in temperature. This is found by subtracting the temperature of the temperature

of the room the chemicals were kept in from the lowest temperature reached in the reaction.

Actual Specific Heat−Experimental Specific HeatActual Specific Heat

∗¿100.0=Correction Factor

Shown in figure 2 below is sample work of the correction factor formula.

Urea Sample Equation

Actual Specific Heat−Experimental Specific HeatActual Specific Heat


333.11−43.58333.11 *100.0=Correction Factor

86.92% = Correction Factor

Potassium Chloride Equation


Actual Specific Heat−Experimental S pecific HeatActual Specific Heat




Ammonium Nitrate Equation

Actual Specific Heat−Experimental S pecific HeatActual Specific Heat




Figure 2. Correction Factor Formula.

The formula used to find the correction factor is found above. This formula is used when

comparing an existing value to an experimental value. By knowing an established value it can

then be compared to the value that was found to find the percent error the statistics holds, to

overall establish a correction factor. The established value for this experiment would be the

researched specific heat for each of the three chemicals. The experimental value would be the

experimental specific heat that was found and averaged out between the trials for each chemical.

The correction factor is used to “fix” the data. What this means is when an established value is

found when researched and the data from the experiment do not match, it can be “fixed” . By

using this formula a proper way, a percentage is found. This percentage or correction factor can

be multiplied by the initial values found to “fix” the data, without redoing trials.

Central Limit theorem justified the normality of the data. Thirty trials were conducted for each


chemical therefore falling under the category for the Central Limit Theorem. This claims if n

represents the number of trials, if n>30, the mean of all samples from the same population will

be approximately equal to the mean of the population. This theorem justifies the normality of the

data for this experiment.

Appendix C: Using a LabProMaterials:

Vernier LabQuestTemperature ProbeFlash Drive

Procedure:

1. Turn on LabQuest and plug in to ensure the device does not run out of battery during trials.

2. Using the stylus, tap the Mode button on the right side of the screen and ensure the mode is

llllset to Time Based.

3. In the Timing section located underneath the Mode selection, make sure to set Length to 180

jgseconds, the Interval to 0.5, and the Rate to 360.

4. Insert the Temperature Probe attachment into the top of the LabQuest.

5. Insert the flash drive you wish to save your data on into the top of the LabQuest.

5. To begin recording data, press the button with a green arrow on it located in the bottom right

jklcorner.

6. If you preset Length correctly, the LabQuest will automatically cease recording data at 180

jjjseconds.

7. To save the data recorded, tap the File button in the top left corner using the stylus. Give the

jjjfile a name, select a desired location to save, and tap Save.


Acknowledgements

As a group we would like to give a special thanks to Mrs.Hilliard for the moral support and

keeping us sane.

We would also like to thank Mr.Supal and Mrs.Dewey for helping us along the course of

research.

We would also like to thank our families for supporting us and standing by our side.

Lastly, we would like to thank our squads for dealing with research talk for numerous months.


Work Cited

Cadrette, and Jaworski. "Temperature Change, Reaction Length, and Specific Heat of

Endothermic Reactions." Temperature Change, Reaction Length, and Specific Heat of

Endothermic Reactions (2009): n. pag. Print.

“The Centeral Limit Theorem.” Wolfram Demonstrations Project (n.d.):1-12.

www.stat.ucla.edu. UCLA.Web. 19 May 2015.

"Collision Theory". Encyclopædia Britannica. Encyclopædia Britannica Online.

Encyclopædia Britannica Inc., 2015. Web. 19 May. 2015

"Enthalpy Changes." SWOT Revision. SWOT Revision, 2011.

Web. 22 Mar. 2015.

“Home Chemistry.” ; Endothermic Reactions. N.p.,

16 Apr/ 2008. Web. 19 May 2015.

“Hot Pack/ Cold Pack.” Nobel.scas.bcit.ca/.

Howard Debeck Elementary School, n.d. Web. 19 May 2015.

“How Do Instant Hot and Cold Packs Work?” www.psa.msu.edu.

Lansing State Journal, 15 Mar. 1995. Web. 19 May 2015.

“Ice and Heat Treatment for Injuries” Patient.co.uk

N.p., n.d. Web. 19 May 2015.

Pasquini, Dr. “Calorimetry Lab” Hartford Physics-.

Tangient LLC, 2015. Web. 22 Mar. 2015

http://www.stat.ucla.edu/


“Sport Injury Statistics.” Stanford Children’s Health. Lucile Packard Children’s Hospital

Stanford, n.d. Web. 19 May 2015.

The Thermodynamics of the Dissolution of Urea (n.d.): 164-70

Www.chemtopics.com. Web. 22 Apr. 2015.

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