Evaluating Which Deicer Is Economically and Environmentally Efficient

Evaluating Which Deicer Is Economically and Environmentally Efficient Abstract The purpose of this experiment was to find the deicers ability to lower the freezing point of water, enthalpy of dissolution, cost per unit, environmental consideration, and later determine the best chemical and economical de-icer through comparison of data found by other groups for opposing de-icers. Potassium chloride was assumed to be one of the best deicers and was the one picked for the experiment. The freezing point depression of KCl, enthalpy of dissociation, van’t hoff factor and calorimeter constant was all tested and determined by running tests. The freezing point depression was tested and registered the first day of the lab with an ice bath. The temperature change was then recorded. The Van’t hoff factor was found by the equation. The whole procedure was repeated again on Day 3 for better results. The Van’t hoff factor was ~1.61. The enthalpy was found on day 2 by the use of a calorimetry. Hot water and cold water were added together in the calorimetry to find the calorimetry constant. Calorimetry constant was found by equations. This trial was run three times. The temperature change of KCl when hot and cold water was presented was recorded and used to calculate the equations. Results gained for the experiment explained that the enthalpy of KCl was 20.96 kJ/mol KCl and endothermic. After comparing results and data with other groups, the conclusion that the best deicer was NaCl overall. KCl was found to be the most expensive, most endothermic, and harmful to the environment. Introduction Minnesota is well known for their below freezing winters. January is the coldest month and temperatures can be seen to be as low as -30 degrees fahrenheit1. Roads are known for quickly gathering lots of ice and making it even more dangerous to travel through it. It can cause slips between wheels and be more difficult to maneuver because of the way the roads are set up which can lead to serious unintentional car accidents. Ice makes it difficult for drivers and also pedestrians. In winter deicer is used to melt ice and break the bond between ice or snow and the pavement to facilitate snow/ice removal. A good deicer is significant for the winter time so finding a high quality deicer is extremely important. Windshield wipers fluid can play an important role in deicing too, because when your windshield is filled with ice instead of getting out the car in below weather to scrape your windshield you could just reduce to ice with the windshield wiper fluid. Ice deposits on windshield affect the vision of the driver therefore creating a safety hazard because it makes it difficult to drive and it’s hard to remove. There is a study being made to add heating to the system of the washer fluid in automobiles, so it can effectively deice the windshield2 . The State of Minnesota Department of Transportation is trying to find the best chemical and economic deicer to use on icy roads to improve safety for drivers. The deicers being considered by the Minnesota Department of Transportation are NaCl, KCl, MgCl2 and CaCl2. There are three traits that make a deicer an effective deicer. First is a deicer that causes a significant decrease in freezing point of water. Second is a deicer that absorbs the most heat when the salt is dissolved which is also known as the enthalpy of dissociation. Lastly is how the deicer acts and affects the environment. In this lab the purpose was to use the given deicer KCl, and test its effectiveness on melting ice on the roads and highways. To accomplish this the freezing point of water when KCl was added to distilled water was measured with an ice bath and using the freezing point depression equation. After finding the freezing point depression the information was taken to find the Van’t hoff factor. Next the enthalpy of dissociation was measured by using a calorimetry and finding the calorimetry constant and plugging that into the enthalpy of dissociation equation. Lastly, the conclusion on how the deicer acts and affects the environment is taken into account. Experimental In this experiment for Day 1 the freezing point of depression was meant to be determined. In order to discover that an ice bath was made. Measured 10 mL of distilled water and added it to a clean test tube. Measured the initial temperature of the distilled water. Measured a 500 mL glass container and filled it with ice. 1.5 grams of rock salt was sprinkled in the ice container. KCl was then measured by different quantities by group members 0.5g, 1.0g, 1.5g and 2.0 grams using a balance. The measured quantities of KCl were added to the distilled water in the test tube. The temperature was then recorded. Solution was mixed together using a glass stir rod. Test tube was then placed into the iced filled container. Two minutes went by then the temperature was recorded. That temperature was known as the freezing point and used in the freezing point depression equation (ΔT=i (mxKf). The Van’t Hoff factor was the i in the equation and was found with the same equation. Van’t Hoff factor was ~1.61. This part of the experiment was repeated again using 1.5 grams of KCl and 1.0 grams of KCl to gain a better result. For Day 2 the enthalpy of dissociation was found. A calorimetry was made using two styrofoam cups, a cardboard lid, some duct tape used to cover properly, and a temperature probe. The heat capacity of the calorimetry was necessary for calculating the enthalpy of KCl, so that was found first. Measured out room temperature water. Recorded temperature. Added ice to the room temperature water and mixed it for 15 seconds. Removed the ice for the water. Measured out hot water. Measured out 23 mL of hot and cold water. Pour both into the calorimetry. Closed and taped lid shut quickly. Used a hole placed in the middle to stir the solution with a glass rod. Recorded the temperature after 2 minutes. Repeated the same steps 2x for a total of 3 trials. Calculated the average to get the calorimetry constant using qcold + qhot + qcal = 0, qcold = Cwater * mcold (Tf - Tcold) ; qhot = Cwater * mhot (Tf - Thot) ; qcal = Ccal (Tf - Tcold). The average calorimetry constant was 12.203 J/C. To find the enthalpy of dissociation 15 mL of distilled water was added to the calorimetry. The temperature was then recorded. Added a given amount of KCl (0.2g, 0.4g, 0.6g, and 0.8g) to the calorimetry. Quickly closed lid. Used glass stirring rod to stir. Measured temperature after 2 minutes. Used equations qreaction  + qwater  + qcal  = 0 , qwater  = CWater  * mwater (Tf - Tinital) , ΔH = qreaction / moles of KCl. Found ΔH. Enthalpy was 20.96 kJ/mol KCl, and endothermic. Results Results for Day 1: Concentration of KCl Freezing point (C°) 0.5g; -2 1.0g; 1.34M -2 1.5g; -3 2.0g; -2 Van’t Hoff factor : ~1.61 Results for Day 2: Finding Calorimetry constant Trial Total Mass (g) Temp (Hot) (C) Temp (cold temp) (C) Final Temp (C) Temp change (C) Temp change (C) (cold) (hot) 1 20g 48 C 20 C 32 C 12 C 2 20g 35 C 20 C 28 C 8C -16 C -7 C 3 20g 60 C 20 C 35 C 15 C -25 C The average calorimeter constant is 12.203 J/C For the enthalpy of adding KCl to water: Trial Mass of KCl (g) Mass of water Temp of water (C) before Final Temperature (C) Temp Change (C) (Final initial) 1 0.2 g 15 g 21.1 20.9 -0.3 2 0.4 g 15 g 22 C 21 C -1 3 0.6 g 15 g 23 C 21 C -2 4 0.8 g 15 g 21 C 18 C -3 C Enthalpy: 20.96 kJ/mol KCl Discussion In the first part of the lab the Van’t Hoff factor was discovered using the equation  ΔT=i (mxKf). For the first trial, the mass of KCl was 1.5 g and the change of freezing point was -3 degrees celsius. The molality of the KCl was 1.5g /74.551g/mol/0.01kg=2.01 M. According to the Van’t Hoff equation ΔT=i (mxKf), 20.5= i(2.01M x 1.61). The same calculation was repeated twice for a total using different grams. By graphing the different freezing points of the KCl concentrations we can determine the Van’t Hoff factor through the slope of the graph. Van’t Hoff factor was ~1.61. The percent error was 24%, because the ideal Van’t Hoff factor is supposed to be 2. For Day 2 of the lab the the heat capacity of calorimeter was discovered at the beginning of the lab with the equations qcold + qhot + qcal = 0, qcold = Cwater * mcold (Tf - Tcold) ; qhot = Cwater * mhot (Tf - Thot) ; qcal = Ccal (Tf - Tcold). The Ccal was found using the equation qcal = Ccal (Tf - Tcold).  he difference between cold and hot was 15 °C and 43 The Ccal was calculated to be 12.203 J/C. T °C. So the qhot = Cwater * mhot (Tf - Thot) , 4.184 J/g°C x 20 x (20-36)= -1338.88 J, and qcold = Cwater * mcold (Tf - Tcold), 4.184 J/g°C x 20 x (32-12)= 1673.6 J. Based on the equation of qhot + qcold+qcal=0, -1338.88, 1673.6+ qcal=0; qcal =334.72. The temperature when cold and hot was mixed was 27 °C. The heat of capacity was 12.203 J/C. qcal = CcalΔT, 334.72= Ccal (35-8).  Therefore the heat capacity of calorimetry is 12.203 J /C. The average of the enthalpy was 16,186 J/mol. The percent error was 6.3%. On day 1 we were unable to get crystallization for 1.5g and 1.0g of KCl for the ice bath. An error could have occurred when one of the lab members was using bottled water including other ingredients rather than just purified water. When the trial was repeated again on Day 3, distilled water was used and crystallization was able to be formed. Another error was on day 1 also the test tube that was placed in the ice bath was left there for five minutes and then the temperature was recorded. It was decided by the group that the time in the ice bath should be limited to two minutes so the exact time of KCl can be recorded effectively. If the experiment were to be repeated, instead of each member in the group performing one trial in the ice bath, all members can perform three separate trials. The percent error for finding the Van’t Hoff was high so performing more trials can improve the accuracy of the data. The data given can be compared and a conclusion can be drawn. Analyzing the cost of KCl as of September 2020 is $2023 per unit, the most expensive deicer in the selection. The enthalpy for KCl is endothermic and in order for a deicer to be good it has to be exothermic, because the heat absorbed through the ice has to be high. KCl is also harmful to the environment. It damages the concrete, known for seeping into groundwater, and is the reason heavy metals are being released. KCl is also known for damaging soil mineral balance. However NaCl is $42 per unit4 , which is much cheaper than KCl. It is also a little endothermic and least harmful to the environment. NaCl causes a decrease in freezing point. The conclusion is that NaCl is the best choice for a deicer. Conclusion For Day 1 of the experiment the Van’t Hoff factor was found which is ~1.61. For Day 2 of the experiment the average enthalpy of dissociation was discovered to be 12.203 J/C. The other thing found on Day 2 of the experiment was the ΔH, which was 1 6,186 J/mol. All this information helped in trying to determine if KCl is the right fit deicer. After reviewing the cost per unit and environmentally affects the conclusion was made that KCl wasn’t the best choice for a deicer. If this work were to be expanded a solubility test would be added to the experiment, because solubility of salts at low temperature is also a major factor, and is used to test environmental concerns. Snow and ice change into liquid state because of solubility, so it's an effective piece to include. Deicers are used by citizens every winter, it's an important use in the act of clearing ice. Less accidents can take place if the effective deicer is used among people. The conclusion has been made that the KCl isn't fit for the best deicer because it’s the most expensive, most endothermic, and harmful to the environment. The best deicer is NaCl because it's cost efficient, a little endothermic, and the least harmful to the environment.