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The Change in Temperature of Urea, Potassium Chloride, and Ammonium Nitrate When Reacting With WaterAmanda Conlon, Sara Nevedal, and Paige RedlinMacomb Mathematics Science and Technology CenterChemistry Mrs. Hillard/ Mrs. Dewey/ Mr. SupalMay 21, 2015Table of ContentsIntroduction…………………………………………………………………...…………………...1Review of Literature………..……………………………………………………………………..4Problem Statement...……………………………………………………………………………..10Experimental Design…………………………………………………………………………..…11Data and Observations………………………………………………………………………..….13Data Analysis and Interpretations………………………………………………………………..20Conclusion……………………………………………………………………………………….28Appendix A………………………………………………………………………………………34Appendix B………………………………………………………………………………………35Appendix C………………………………………………………………………………………38Acknowledgements………………………………………………………………………………39Work Cited……………………………………………………………………………………….40IntroductionIt’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 LiteratureNo 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 DiagramsFigure 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 SystemFigure 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 Ch?telier'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 StatementProblem: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:Calorimeter325 g of Ammonium Nitrate, NH4NO3TI-nspire calculator 250 g of Urea, CH4N2O150 mL Graduated Cylinder300 g of Potassium Chloride, KClLabQuestScale (To the nearest 0.0001g)(90) Weigh BoatsLabQuest 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 CalorimeterFigure 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 ObservationsData:Table 1 Urea Reaction Temperature Changes TrialInitial Temperature of Water(°C)Lowest Temperature Reached(°C)Change in Temperature (°C)120.517.6-2.9220.617.4-3.2320.317.2-3.1420.117.2-2.9520.817.9-2.9622.920.4-2.5721.718.8-2.9822.019.1-2.9921.118.1-3.01020.617.7-2.91123.620.6-3.01221.419.3-2.11321.118.0-3.11423.520.2-2.31522.719.1-3.61622.018.8-4.21721.518.4-3.11821.217.6-3.61920.417.9-2.52020.918.5-2.42120.918.4-2.52221.118.5-2.62320.017.8-2.22420.118.5-1.62522.020.0-2.02621.818.8-3.02721.518.2-3.32821.117.5-3.62920.917.6-3.33020.818.3-2.5Average21.318.2-3.1Table 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 ChangeTrialInitial Temperature of Water(°C)Lowest Temperature Reached(°C)Change in Temperature (°C)119.816.1-3.7220.416.7-3.7320.216.6-3.6420.917.3-3.6521.017.3-3.7620.917.2-3.7721.117.5-3.6821.117.4-3.7919.916.4-3.51019.816.3-3.51119.816.2-3.61220.216.3-3.91319.916.3-3.61419.516.2-3.31520.917.5-3.41620.616.8-3.81719.816.5-3.31819.316.1-3.21919.816.3-3.52019.716.4-3.32120.016.8-3.22219.916.5-3.42319.416.3-3.22419.816.5-3.32520.216.7-3.5TrialInitial Temperature of Water(°C)Lowest Temperature Reached(°C)Change in Temperature (°C)2620.116.8-3.32720.216.7-3.52819.716.5-3.22919.516.2-3.33020.716.9-3.8Average20.116.6-3.5Table 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 ChangeTrialInitial Temperature of Water(°C)Lowest Temperature Reached (°C)Change in Temperature (°C)120.915.4-5.5220.116.6-3.5321.617.8-3.8422.317.4-4.9522.116.9-5.2621.918.5-3.4722.617.9-4.7822.117.5-4.6921.717.4-4.31022.117.4-4.71124.219.8-4.41225.721.3-4.41324.820.5-4.31423.319.6-3.71522.618.1-4.5TrialInitial Temperature of Water(°C)Lowest Temperature Reached(°C)Change in Temperature (°C)1622.018.8-4.21721.517.4-4.11821.216.6-4.61921.417.9-3.52020.917.3-3.62121.917.9-4.02221.116.5-4.62320.016.8-3.22420.116.4-3.52522.018.1-3.92621.516.9-4.62720.616.5-4.12823.219.3-3.92922.118.2-3.93021.817.6-4.2Average22.017.8-4.2Table 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 ObservationsTrialObservations1Researcher C measured, Researcher B stirred. Water turned cloudy immediately, then at about 100 seconds turned clear again.2Researcher 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.3Researcher C measured, Researcher B stirred. Still small particles after 100 seconds.4Researcher C measured, Researcher B stirred. Cloudiest water yet.5Researcher C measured, Researcher B stirred. Cloudy to clear at 120 seconds.6Researcher C measured, Researcher B stirred. Some air bubbles while stirringTrialObservations7Researcher C measured, Researcher B stirred. Consistently clear, temperature went up 1 degree at first, then dropped about 30 seconds in.8Researcher C measured, Researcher B stirred. Water cleared earlier than 100 seconds.9Researcher C measured, Researcher B stirred. Spilled a little water around rim, cloudy for longer than 100 seconds.10Researcher C measured, Researcher B stirred. Turned clear quicker than 100 seconds.11-12Researcher C measured, Researcher B stirred. Urea was more powdery since it was bottom of tub, never very cloudy, cleared up faster than 100 seconds.13Researcher C measured, Researcher B stirred. Urea was more powdery since it was bottom of tub, very cloudy, cleared up after 100 seconds14Researcher 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.15Researcher C measured, Researcher A stirred. Urea was clumpy before starting reaction. Turned clear at 100 seconds.16- 17Researcher C measured, Researcher A stirred. Turned clear before 100 seconds.18Researcher C measured, Researcher A stirred. Still slightly cloudy after 100 seconds.19-21Researcher C measured, Researcher A stirred. Solution was overall very cloudy.22Researcher 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.23Researcher B measured, Researcher A stirred. Solution turned clear after 100 seconds.24Researcher B measured, Researcher A stirred. Solution started turning clear before 100 seconds.25Researcher B measured, Researcher B stirred. Turned clear before 100 seconds.26Researcher B measured, Researcher B stirred. Slightly clear around 60 seconds, completely clear around 100 seconds.27Researcher B measured, Researcher A stirred. Solution started turning clear before 100 seconds.28Researcher 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. 30Researcher 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 ObservationsTrialObservations1Researcher A weighed, Researcher C stirred. Only slightly cloudy. 2Researcher A weighed, Researcher C stirred. Slightly cloudy, some bubbles around sides.3-7Researcher A weighed, Researcher C stirred. Clear at 70 seconds.8Researcher B weighed Researcher C stirred, large chunks in powder.9Researcher B weighed, Researcher C stirred, clear at 70 seconds.10Researcher B weighed, Researcher C stirred, water splashed.11Researcher A weighed, Researcher C stirred, clear at 70 seconds, some chloride got onto top part of calorimeter.12Researcher A weighed, Researcher C stirred, some chunks in the powder.13-14Researcher C weighed, Researcher A stirred, clear at 70 seconds.15Researcher A weighed, Researcher B stirred. Clear at 75 seconds.16Researcher A weighed, Researcher B stirred. Clear at 80 seconds.17Researcher A weighed, Researcher B stirred. Clear at 70 seconds.18Researcher A weighed, Researcher B stirred. Clear at 75, water splashed out. 19Researcher A weighed, Researcher B stirred. Clear at 75.20Researcher A weighed, Researcher B stirred. Clear at 80, some particles left at bottom. 21Researcher A weighed, Researcher B stirred. Clear at 85 seconds. 22Researcher C weighed, Researcher B stirred, clear at 75 seconds. 23Researcher C weighed, Researcher B stirred. Clear at 80 seconds. 24-27Researcher C weighed, Researcher B stirred, clear at 80 seconds. 28Researcher C weighed, Researcher A stirred, clear at 70 but some particles remained.29Researcher C weighed, Researcher A stirred, clear at 75 seconds.30Researcher 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 ObservationsTrialObservations1-3Researcher C stirred, Researcher A weighed. Very cloudy. Cleared up around 150 seconds. 4Researcher C stirred, Researcher A weighed. Some small particles remained after 100 seconds.5-7Researcher C stirred, Researcher A weighed. Very cloudy. Cleared up around 150 seconds. 8Researcher C stirred, Researcher A weighed. Very cloudy. Cleared up around 115 seconds. 9Researcher B stirred, Researcher A weighed. Clear at 120 seconds.TrialObservations10-15Researcher B stirred, Researcher A weighed. Cloudy, clear at 110 seconds.16-30Researcher 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 ReactionsThe 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 ReactionThe 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 ReactionsThe 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 ReactionsFigure 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 ReactionsFigure 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 NitrateInitial Specific Heat(J/g°C)Correction Factor (%)Specific Heat with Correction Factor (J/g°C)Actual Specific Heat (J/g°C)Urea4.4*10186.923.8*1035.6*103Potassium Chloride3.9*10191.023.6*1035.9*103Ammonium Nitrate4.3*10188.113.8*1034.6*103Table 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. ?ConclusionWhen 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: RandomizationMaterials:Ti-Nspire CX calculatorProcedure:Turn on calculator and open a calculator page.Press the menu button.Scroll down and choose the “5: Probability” by pressing enter.Scroll down and choose the “4: Random” option by pressing enter.Select the “2: Integer” function and press enter.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.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. ?Press enter. A number will appear, representing a trial number. Continue pressing enter until all trials are accounted for. If a number repeats, do not take it into account.Repeat steps 1-9 for each run. Appendix B: Descriptive Data and Analysis EquationsThroughout 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 ΔTShown in figure 1 below is sample work of the specific heat formula.Urea Trials Sample EquationSM ΔT=SM ΔT4.184(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)=S4.4*101 kj/mol=SPotassium Chloride Trials Sample EquationSM ΔT=SM ΔT4.184(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)=S3.9*101 kj/mol=SAmmonium Nitrate Trials Sample EquationSM ΔT=SM ΔT4.184(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) =S4.3*101 kj/mol=SFigure 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 FactorShown in figure 2 below is sample work of the correction factor formula.Urea Sample Equation Actual Specific Heat-Experimental Specific HeatActual Specific Heat*100.0=Correction Factor333.11-43.58333.11*100.0=Correction Factor86.92% = Correction FactorPotassium Chloride EquationActual Specific Heat-Experimental Specific HeatActual Specific Heat*100.0=Correction Factor436.68-39.23436.68*100.0=Correction Factor91.02% = Correction FactorAmmonium Nitrate Equation Actual Specific Heat-Experimental Specific HeatActual Specific Heat*100.0=Correction Factor365.55608-43.4492365.55608*100.0=Correction Factor88.11% = Correction FactorFigure 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 DriveProcedure: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.AcknowledgementsAs 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 CitedCadrette, 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.stat.ucla.edu. UCLA.Web. 19 May 2015."Collision Theory".?Encyclop?dia Britannica. Encyclop?dia Britannica Online.Encyclop?dia 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?” psa.msu.edu.Lansing State Journal, 15 Mar. 1995. Web. 19 May 2015.“Ice and Heat Treatment for Injuries” Patient.co.ukN.p., n.d. Web. 19 May 2015.Pasquini, Dr. “Calorimetry Lab” Hartford Physics-.Tangient LLC, 2015. Web. 22 Mar. 2015“Sport Injury Statistics.” Stanford Children’s Health. Lucile Packard Children’s HospitalStanford, n.d. Web. 19 May 2015. The Thermodynamics of the Dissolution of Urea (n.d.): 164-70Www.. Web. 22 Apr. 2015. ................
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