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Deicer Lab: Determining the Effectiveness of Calcium Chloride (CaCl2) as a Deicer in Comparison to Other Salts i Value, Enthalpy of Dissolution, and Environmental Factors
Emily Young March 3, 2014
Chemistry 1062, Professor: Michelle Driessen, TA: Victoria Szlag University of Minnesota Twin Cities
Abstract
The effectiveness of calcium chloride (CaCl2) as a deicer in comparison to other salts was evaluated based on three criteria its Van’t Hoff factor, enthalpy of dissolution, and impact on the environment. The Van’t Hoff factor (i) was experimentally determined through freezing point measurements of 4 calcium chloride solutions with known molalities. A calorimeter was built in order to determine the enthalpy of dissolution through recording the amount of heat released by the dissolution of calcium chloride. With experiments showing a comparatively high i value and highly exothermic enthalpy of dissolution, calcium chloride was found to be a chemically ideal deicer. These results, along with the cost and environmental impacts of calcium chloride in comparison to sodium chloride (the status quo deicer) and other deicers, make it a favorable deicing alternative.
Introduction Deicers are essential towards improving traffic and making safe road conditions during the winter “up to 12 million tons (of sodium chloride, the most commonly used deicer) are used in the United States each year” 1. However, deicers can potentially harm the environment in several ways they can cause plant dehydration, the runoff into streams can cause excessive chloride levels toxic to aquatic life, and the sinking of dense salty water can impair water circulation, just to name a few. For these reasons, sodium chloride and alternative deicers have frequently been tested to determine which deicer is best, considering chemical effectiveness, impact on the environment, and cost. In this study, the effectiveness of calcium chloride (CaCl2) as a deicer in comparison to other salts is determined, including sodium chloride (NaCl), magnesium chloride
(MgCl2), and potassium chloride (KCl) using experimentally determined chemical as well as economic criteria. Chemical deicers work by preventing the binding of water molecules above the freezing point of water, 0 °C. In this paper, several factors of calcium chloride will be discussed including it’s i (Van’t Hoff factor) value, enthalpy of dissolution, and environmental impacts. These factors conclusively prove calcium chloride and magnesium chloride to be superior to other deicers, namely sodium chloride. Prediction Rearranging the expression for freezing point depression for the Van’t Hoff factor gives
i= ∆TKfm (1)
where is the ∆T is the temperature difference between the freezing point of pure water (0°C) and the freezing point of solution in degrees Celsius, Kf is the freezing point depression constant (1.86°C kg/mol for water), and m is the molality of the solution (moles of solute per kilogram of solvent). The enthalpy of dissolution is given by
ΔHrxn= qrxn# of moles (2)
where qrxn is the heat of reaction. Rearranging for the heat of reaction in the dissolution of calcium chloride in water using the first law of thermodynamics (the law of conservation, qtotal=0) gives
qrxn= (qsol + qcal) (3) The heat input of the reaction is given by
qsol= cm∆T (4) where c is the specific heat capacity of the solution (4.184 J/g °C for water), m is the mass of the solution, and is the ∆T is the temperature change of the reaction in degrees Celsius. The heat input of the calorimeter is given by
qcal= ccal∆T (5) where Ccal is the heat capacity of the calorimeter and ∆T is the temperature change in degrees Celsius. Since the heat capacity of the unknown, it can be determined using a reaction between hot and cold water. Using the expression for heat and the first law of thermodynamics and solving for the heat capacity of the calorimeter gives
Ccal=Cw[ ](Tf−Tc)MH(Tf−TH) + Mc(Tf−Tc) (6)
where Ccal is the heat capacity of the calorimeter, Cw is the heat capacity of water (4.184 J/g °C for water), MH is the mass of the hot water, Tf is the final temperature, TH is the initial temperature of the hot water, and Tc is the initial temperature of the cold water (the initial temperature of the cold water and the calorimeter were assumed to be same for this experiment). Rearranging equations #26 for the enthalpy of dissolution and simplifying gives
ΔHrxn=∆TCw[ m](Tf−Tc)MH(Tf−TH) + Mc(Tf−Tc) (7)
Equation 7 predicts the enthalpy of dissolution. Procedure
Determining Van’t Hoff Factor
Figure 16
Figure 1 depicts the mechanism used to determine the Van’t Hoff factor (i). Massed samples of calcium chloride were placed into test tubes of massed water and mixed thoroughly until dissolved. A temperature probe/thermistor was inserted into the sample, which was inserted into an insulated rock and ice salt bath prepared in a 200mL beaker, recording the freezing point with Logger Pro. The average value for the portion of the graph with constant temperature was recorded as the freezing point. The temperature readings were used in equation 1 to determine the Van’t Hoff factor.
Determining Enthalpy of Dissolution
Figure 27
Figure 2 depicts the mechanism used to determine the enthalpy of dissolution. The heat capacity of the styrofoam cup calorimeter was first tested by placing cold water in the styrofoam cup, adding hot water, recording the initial and final (equilibrium) temperatures using a temperature probe and Logger Pro, and plugging the temperature and mass values into equation 6. After the heat capacity was determined, calcium chloride samples were massed and placed into the calorimeter. A temperature probe was immediately inserted and used to stir the solution while recording the temperature using Logger Pro. The temperature of the solution was recorded at its peak. The experimental values were used with equation 4 to determine the enthalpy of dissolution.
Data
Van’t Hoff Factor (i) Freezing point data from calcium chloride solutions:
Data from determining Van’t Hoff factor (i): molar mass of calcium chloride = 111.7g
Trial Mass of water (g)
Mass of calcium chloride (g)
Freezing Point (°C)
molality (moles calcium chloride/kg
calculated Van’t Hoff factor
water)
1 9.320±.001
0.104±.001
0.48
0.0999±.002
2.74± .177
2 9.278±.001
0.303±.001
1.50
0.2924±.002
2.81± .177
3 10.137±.001
0.509±.001
2.4
0.4495±.002
2.91± .177
4 10.365±.001
0.704±.001
3.0
0.6081±.002
2.61± .177
Average i Value: 2.78± .177
Enthalpy of Dissolution Calorimeter Data:
Trial Temp (hot) (°C)
Temp (cold) (°C)
Temp (final) (°C)
Mass (hot) (g)
Mass (cold) (g)
Calculated Ccal value
1 65.4±.01
3.7±.01
34.7±.01
56.664±.001
52.422±.001
15.45±3.3
2 60.9 ±.01
3.7±.01
26.4 ±.01
47.991±.001
66.831±.001
22.4±3.3
Average Ccal value = 18.925±3.3 J/°C Enthalpy of Dissolution Data:
Trial Mass water (g)
Mass calcium chloride (g)
Water temp (°C)
Solution temp (°C)
Calculated enthalpy of dissolution (kJ/mol)
1 50.701±.001
0.214±.001
20.3±.01
20.8±.01
60.30 ±15.1
2 76.930±.001
0.227±.001
20.6±.01
21.0±.01
67.08±15.1
3 76.260±.001 0.254±.001 20.7±.01 21.3±.01 89.18±15.1
Average enthalpy of dissolution: 72.19±15.1kJ/mol
Environmental Impacts
Environmental Impact Data:
NaCl CaCl2 MgCl2 KCl
soils Na can bind to soil particles, breakdown soil structure, and decrease permeability. Cl can form complexes with heavy metals in the soil, releasing them into the environment.
Cl may form complexes with heavy metals, releasing them into the environment. Ca improves soil structure.
Cl may form complexes with heavy metals, releasing them into the environment. Mg improves soil structure.
K can exchange with heavy metals in soil, potentially releasing them into the environment.
vegetation Traffic spray may cause leaf scorch and browning. Uptake of NaCl may cause osmotic stress, and nutrient imbalance.
May cause osmotic stress and leaf scorch. Ca is an important macronutrient for plant growth.
May cause osmotic stress and leaf scorch. Mg is an important element in plant physiology.
Similar to NaCl (osmotic stress and leaf scorch).
groundwater Elevated levels of Cl near roadways.
Similar to NaCl, cation exchange of Ca may increase potential for metal contamination.
Similar to NaCl and CaCl2.
Similar to NaCl, K can release heavy metals from soil into groundwater.
surface water Excessive chloride loading possible. Can cause saline stratification that may lead to anoxia in lake bottoms.
Similar to NaCl.
Similar to NaCl.
Similar to NaCl.
aquatic life No effect on large bodies of water, but harmful concentration of Cl possible in streams.
Similar to NaCl.
Similar to NaCl.
Similar to NaCl.
human use Mild irritant to skin/eyes.
Similar to NaCl.
Similar to NaCl.
Similar to NaCl.
Average Cost (in $ per metric ton) 2,5: calcium chloride: $120315 sodium chloride: $51.70 magnesium chloride: $111122.32 potassium chloride: $323
Analysis Freezing point:
After each of the solutions experienced instantaneous freezing, the temperatures recorded on logger pro jumped and leveled off at a constant temperature this constant represents the freezing point of the solution. The temperature of the solution when it froze instantly was not recorded as the freezing point, because the solution was supercooled at that point. The right part of the graph for each trial shows where the temperature of the solution reached a near constant temperature. The mean of this constant was recorded as the freezing point 0.48 °C, 1.50 °C, 2.4 °C, and3.0 °C, respectively.
The trend observed here is that, as the mass of calcium chloride was increased in each solution, the freezing point decreased. This can be explained by colligative properties the addition of ions to a solvent decreases the vapor pressure of the liquid solvent in this case, calcium cations and chloride anions decreased the vapor pressure of liquid water. Raoult’s law (Psolvent = XsolventPosolvent) explains this behavior the vapor pressure of the solvent is equal to the mol fraction of the solvent multiplied by the vapor pressure of the pure solvent, so the vapor pressure of the solvent above the solution decreases as more solute is added and the fraction of solvent is decreased. This occurs because, as solute is added, the ions block solvent particles from escaping. In this case, as calcium chloride was added to water, calcium cations and chlorine anions blocked water from escaping, lowering the vapor pressure. This resulted in a lowering of the freezing point the freezing point occurs when the vapor pressure of the liquid is equal to the vapor pressure of the solid. The normal equilibrium for water where this occurs (0 °C) is disrupted when a solute is added. More specifically, the addition of calcium and chlorine ions diluted the water so that less water was available to freeze, lowering the point at which the vapor pressure of liquid water was equal to the vapor pressure of solid water resulting in a lower freezing point.
Determining i value:
First, the mechanism through which calcium chloride decreases the freezing point of water will be discussed. Simply put, adding salts prevent water molecules from binding to form ice. Colligative properties were used to explain the phenomena of freezing point depression with the addition of solute, but it can also be explained through entropy and enthalpy. Salts work to decrease the freezing point of water by contributing to a battle between entropy and enthalpy. The universe tends towards higher entropy while favoring reactions that decrease enthalpy. For pure water, the phase change from liquid to solid occurs at the point where the influence of enthalpy overcomes the influence of entropy at 0°C. This balance is shifted with the addition of solute while solute has no effect on water in it’s solid form, it increases both enthalpy and entropy of water in it’s liquid form but predominantly entropy. Therefore, because the addition of calcium chloride raised the entropy, or disorder, of water in it’s liquid form, the transition to it’s solid form became less favorable, and a lower enthalpy reduced temperature was required in order for freezing to occur. Secondly, the i values will be discussed the i value is a measure of the impact of a solute on colligative properties, in this case freezing point depression. The i value depends not on the identity of the substance but on it’s concentration, meaning a compound that separates into more ions (calcium chloride and magnesium chloride) should have a higher i value the results were consistent with this. The average i value for calcium chloride was 2.78± .177, relatively close to the ideal value of 3, and higher than the i values for
any of the other deicers. These i values can be interpreted in a few ways on the one hand, a higher i value means more dissociated ions and therefore, a stronger environmental impact per mol of deicing compound (something that will be discussed later). On the other hand, this also means that a smaller amount of compound will accomplish the same job which minimizes environmental impact. With comparatively minimal environmental impacts (again, discussed later), the high i value of calcium chloride makes it the best deicer in this category. The low freezing point of calcium chloride solutions makes calcium chloride a much more effective deicer than the widely used sodium chloride, having the capability to melt ice well below its freezing point. With just 4 trials of calcium chloride solutions, the data acquired is limited. More trials would have given a more accurate estimate for the i value of calcium chloride. Room for error must be accounted for as well. Systematic error exists in addition to the potential human error in the temperature readings the solute may not have been completely dissolved, which would result in a higher freezing point and less than ideal Van’t Hoff factor. Other solutes may have gotten into the test tubes and contaminated the solution, resulting in a potentially lower freezing point that doesn’t reflect the actual freezing point depression capability of calcium chloride and would have increased the i value.
Calorimeter:
The average Ccal value for the calorimeter built was calculated to be 18.925±3.3 J/°C. This value was used to calculate the enthalpy of dissolution for calcium chloride. Using a styrofoam cup calorimeter allowed for retention of heat in order to achieve more accurate enthalpy of dissolution data. However, with only 2 trials, this data is limited a more extensive experiment would have included more trials to get a more accurate Ccal value, or perhaps built more calorimeters for a more extensive enthalpy of dissolution test. Room for error must also be accounted for heat loss during transport of water from beaker to calorimeter.
Enthalpy of Dissolution:
The average enthalpy of dissolution for calcium chloride over 3 trials was 72.19±15.1kJ/mol. This indicates that the dissolution of calcium chloride is a highly exothermic reaction, meaning it releases heat when it reacts with water making it an effective deicer. Nonexothermic salts (sodium chloride, for example) may work down to certain temperatures the sun can provide thermal energy for it to melt ice, and it can also work through cars driving over the salt and grinding it into the ice, providing a barrier that prevents sliding. However, salts with a highly exothermic dissolution reaction have a much more powerful ability to melt ice, particularly at lower temperatures they release heat upon dissolving in water, adding thermal energy to the system. The enthalpy of dissolution for calcium chloride was more exothermic than potassium chloride or sodium
chloride, but not quite as exothermic as magnesium chloride, placing it second in this category 5. More trials could have been done to reach a more accurate average enthalpy of dissolution value with only 3 trials, these values are limited. Room for error must also be accounted for incomplete dissolving of calcium chloride would have resulted in a lower temperature recording, and thus, a calculation that shows a less exothermic reaction than reality. Heat loss and the limitations of the calorimeter may also have resulted in an inaccurate enthalpy of dissolution calculation. Human error in reading temperature is also probable, as well as potential loss of calcium chloride between measurement of mass and mixing into solution that would have resulted in a heat of enthalpy calculation that was less exothermic than it should be for the recorded mass.
Environmental impacts: The effects of calcium chloride on the environment are comparable to that of potassium chloride and magnesium chloride, and less harmful than sodium chloride. Both calcium chloride and magnesium chloride contain ions that improve soil structure, a benefit that sodium chloride lacks. This makes calcium chloride and magnesium chloride good alternatives to sodium chloride, the most commonly used deicer, when considering environmental impacts.
Cost:
Calcium chloride, on average, costs more than sodium chloride or magnesium chloride, but is less expensive than potassium chloride. Sodium chloride is the least expensive deicer, but this doesn’t take into account environmental costs, which are hard to quantify but certainly outweigh the initial cost of an alternative deicer. Therefore, it is difficult to say which deicer is best economically.
Conclusion With an average experimentally determined i value of 2.78± .177 and enthalpy of dissolution of 72.19±15.1kJ/mol, the dissolution of calcium chloride is highly exothermic, meaning it releases heat making it chemically a more effective deicer in comparison to sodium chloride and potassium chloride, but not as effective as magnesium chloride 5. Sodium chloride is largely ineffective as a deicer below 10 °C, while calcium chloride and magnesium chloride work well below this temperature 5. Both calcium chloride and magnesium chloride have higher i values than sodium chloride and therefore require a smaller amount to melt the same amount of ice. With environmental impacts comparable to to magnesium chloride and potassium chloride and
less harmful than sodium chloride, calcium chloride is also a top contender in terms of environmental concern. Lastly, calcium chloride costs more than sodium chloride or magnesium chloride but less than potassium chloride per metric ton, making it not the best choice economically but still a viable one, especially when considering the environmental costs that sodium chloride cause, which aren’t included in the price listed above. Overall, the results from the experiments suggest that both magnesium chloride and calcium chloride would make good alternatives to sodium chloride. With sodium chloride currently being the most widely used deicer by a large margin, our experiments show that a change is in order. The use of calcium chloride or magnesium chloride in place of sodium chloride would reduce the negative impact of deicing on the environment, as well as provide more effective deicing in colder temperatures. References: 1). Carder Ledyard & Milburn LLP. Environmental Impact of Road Salt and Deicers. http://www.clm.com/publication.cfm?ID=321 (accessed 2/24/14) 2). Michigan. Effects of Deicing Materials on Natural Resources, Vehicles, and Highway Infrastructure. http://www.michigan.gov/documents/ch3deice_51440_7.pdf (accessed 2/24/14) 3). Adirondack Watershed Institute. Review of Effects and Costs of Road DeIcing with Recommendations for Winter Road Management in the Adirondack Park. http://www.adkaction.org/files/public/Full_Study_Salt.pdf (accessed 2/24/14) 4). Silberberg, M. Chemistry: The Molecular Nature of Matter and Change, 6th ed.; McGrawHill: United States, 2012. 5). Szlag, V. Deicer 230, University of Minnesota: United States, Februrary 28, 2014. 6). Indiana University. http://www.iupui.edu/~cletcrse/380/ch3suppos_files/image002.jpg (accessed 2/24/14) 7). University of Texas. http://ch301.cm.utexas.edu/images/coffeecup.png (accessed 2/24/14)
