Abstract
A common road salt, magnesium chloride, was analyzed in the lab to test its effectiveness as a road salt. The Van’t Hoff factor of the salt and the enthalpy of dissolution when dissolved in pure water were experimentally tested to evaluate its many characteristics as a deicer. To determine the Van’t Hoff factor, the difference of temperatures of freezing water and a solution of magnesium chloride was determined. A Van’t Hoff factor average of 2. 88 was found, which is close to the ideal Van’t Hoff factor value of 3. The enthalpy of dissolution was found using calorimetry. Various amounts of magnesium chloride were added to water finding the heat of reaction for each trial.
Depending on how many moles of magnesium chloride were present in each solution, the enthalpy of dissolution was then calculated which gave an average of -151. 45 kJ/mol making the reaction exothermic. Magnesium chloride was compared to three other road salts to determine which of the road salts was the best for Minnesota’s winter roads. Due to the ability to lower the freezing point of water, the substantial value of enthalpy change, the least cost effective, and non-harmful effects to the environment it makes the best chloride deicer out of potassium chloride, sodium chloride, and calcium chloride.
Introduction
Minnesota roads are known to become increasingly dangerous once temperatures reach below 0°C. To make the roads less of a hazard in the winters, the State of Minnesota Department of Transportation has adapted to coat the roads with salts to deice the roads. Deicers are used to eliminate or reduce the amount of ice present in icy and snowy conditions. This experiment will test the effectiveness of a specific salt, magnesium chloride (MgCl2), as a deicer for Minnesota roads.
The freezing point depression of magnesium chloride in water and the heat of enthalpy of magnesium chloride in water will be tested to help in determining how effective of a deicer magnesium chloride is. The freezing point depression will be calculated by monitoring the temperature drop of different saturations of magnesium chloride in ice baths. The heat of enthalpy will be calculated by adding different masses of magnesium chloride to a calorimeter and observing the change of heat for each trial which will give information on whether the reaction of magnesium chloride and water is an exothermic or endothermic reaction.
These results along with cost per kg and environmental factors will be compared to the other salts in the lab, NaCl, KCl, and CaCl2, to find which of the salts is the best for deicing Minnesota’s winter roads.
Experimental
Experiment 1
Four trials of ice baths were carried out using various amounts of magnesium chloride in each trial. All trials were performed filling up a 400 mL beaker about three fourths full of ice and adding 10. 0 mL of deionized water to each 400 mL beaker of ice. Using a LoggerPro temperature probe, the initial temperature of each the ice baths were recorded to test accuracy of the temperature probe.
The following masses: 2. 002 g, 2. 003 g (a second time) 4. 003 g, 4. 003 g (a second time), and 5. 002 g of magnesium chloride were then poured into each beaker of ice, each mass being used for one trial each. The LoggerPro temperature probe was used to stir the solution to ensure evenly dispersion of magnesium chloride. Once the freezing point reached a constant temperature by monitoring the graph on the LoggerPro program, the value at constant determined the freezing point of depression of the new solution by taking the difference between the experimental freezing point of water (0°C), and the new freezing point of the solution.
Also, the percent of ice remaining in the beaker was observed after the final temperature was reached. The additional ice that melted was added to the initial mass of water to obtain the final weight of water in the solution to calculate molality: molality m= (grams MgCl2? 95. 21gmol of MgCl2)10mL H2O? 1g H2O1 mL H2O? 1 kg H2O1000 g H2O+grams of ice? % of ice melted? 1 kg ice1000 g ice To obtain the average Van’t Hoff factor from the data, plot the change of temperature vs. molality and find the line of best fit. The slope represents the average Van’t Hoff factor multiplied by the freezing point constant of water, (i)(Kf).
Then calculate each trial’s Van’t Hoff factor by using Equation 2. This originated from the freezing point depression equation: i = ? T/((Kf)(m)) (Equation 2) The freezing point constant (Kf) of water is 1. 86 °C m-1. Each mass amount and Van’t Hoff factor was calculated then analyzed in a table.
Experiment 2
Experiment 2 focused on finding the enthalpy of solution of magnesium chloride. Testing the enthalpy of solution started with measuring out 10 mL of deionized water in a graduated cylinder for three separate trials, each trial having a different mass of magnesium chloride.
The water was then poured into a well of a Styrofoam calorimeter then the initial temperature of water was taken using a temperature probe and the LoggerPro programming. A measured amount of magnesium chloride was placed in the same well as the water in the calorimeter, immediately covering the well and stirring the solution with the probe. With the temperature probe in the solution, the change in temperature of the solution was monitored every 30 seconds and the greatest change observed in the LoggerPro graph would then be used for change in temperature.
The following equations were used to find the change in heat of the magnesium chloride water solution: C solution = (W1)(C solid) + (W2)(C water) Equation 3 (Dimoplon Equation1) Q solution = (m solution)(C solution)( ? T) Equation 4 Q cal = (C cal)( ? T) Equation 5 Q rxn = -Q cal – Q solution Equation 6 H = (Q rxn)/(mol MgCl2) Equation 7 (Enthalpy Dissolution) The calorimeter used in this experiment was calculated to have a specific heat capacity value of 15. 4 J/°C.
The specific heat capacity of water is 4. 184 J/(kg*°C). The heat of solution was determined by plotting Q rxn of all trials vs. the number of moles of magnesium chloride. The slope of the line of best fit is equal to the heat of solution of magnesium chloride.
Experiment 3
The freezing point depression experiment was retested in hopes of reaching more accurate results with changes from experiment 1. 50 mL of deionized water was measured out in a 100 mL graduated cylinder three times and each amount poured into three separate 50 mL beakers. The following masses: 2. 000 g, 4. 000 g, and 5. 00g of magnesium chloride were then poured into the 50 mL beakers, each mass having their own labeled beaker.
To ensure that all of the solute dissolves in the water, the solution was stirred using a stirring rod. Fill a 200 mL beaker three fourths full of ice and add 25 mL of deionized water. Using a temperature probe monitor the temperature change as rock salt is added to the ice bath until the temperature reaches -10. 0o C. Once the desired temperature is reached, one of the 50 mL beakers containing the magnesium chloride solution was submerged in the ice bath. The ice bath was continually monitored, keeping it at the same temperature.
The lowest temperature was then recorded, which was constant for at least 30 seconds. Each beaker of solution was submerged in the ice bath. Each temperature change was recorded and molality was calculated using the same process of plotting the points of each trial from experiment 1.
Using the slope of the trend line, it can be concluded that 151. 45 kJ of heat is released for every increase of 1 molality of MgCl2. The negative slope also shows that the reaction is exothermic. Discussion The theoretical Van’t Hoff of MgCl2 is 3, because MgCl2 breaks into three ions when dissolved completely in water. The results from Experiment 1 showed a Van’t Hoff of 1. 45%, a 52% error. Experiment 1 had many errors involved, which can account for the inconsistent results in that experiment. Due to the inconsistency, the new Van’t Hoff factor calculated in Experiment 3 was 2. 88, with a 4% error.
Although the Van’t Hoff factor numbers were not satisfactory, the magnesium chloride clearly lowered the freezing point of water, shown in graphs and tables 1 and 2. As the molality of the solution increased, the increases concentration of magnesium chloride lowered the freezing point, but only to a certain molality did this seem true. As the molality increased in small increments, the change in temperature increased in a positive relationship. In both Experiment 1 and 3, the last trial that consisted the most grams of magnesium chloride in the solution showed the opposite trend of increase concentration equals lower freezing point.
Experiment 1’s data may have differed more compared to Experiment 3 because of an inconsistency of where the “bend” in the freezing point graph location was placed. Another factor would be that in Experiment 1 mixed the ice with the magnesium chloride solution directly. This made finding the final mass of the solvent a difficult task because as the magnesium chloride is being stirred into the ice bath, the ice melted at a rate in which could then only be accounted for by visually estimating how much ice was left in the beaker.
By visually estimating the amount of ice melted this significantly skewed our results in calculating the molality of the solution. Taking the errors from Experiment 1 into account, rock salt was used in Experiment 3 in hopes to eliminate error in the mass of the solvent. Although this helped in making our ice bath cold, the rock salt only kept in the solution for a short period of time before the ice bath began to warm up, especially towards the end of the experiment when the molality of magnesium chloride in the solution increased.
As more magnesium chloride was added to the 50 mL beaker, the warmer the solution became and the more difficult it was to for the solution to reach its freezing point. Another factor to take into account between Experiment 1 and 3 are the number of trials in each. In both days, the molality of each trial could have been closer in value instead of so far apart to really determine a relationship between molality and freezing point depression. From looking at the tables from both experiments the trend of increased molality held true until a certain molality was reached, any value above 0. 0 m seemed to throw off the results.
Theoretically, as the concentration of magnesium chloride in a solution increases, the amount that dissociates will decrease due to the higher concentration on Mg2+ and Cl- ions in the solution. More trials could have supported this theory more accurately. The solubility of magnesium chloride in water is 52. 9 g/mL at 0 degrees Celsius which can also account to the unfitting Van’t Hoff factors when the molality is at its highest because the solutions may have been too saturated2. The results from day two show that magnesium chloride has a negative enthalpy of solution.
This proves that magnesium chloride is exothermic when mixed with water. An exothermic reaction releases heat to its surroundings. As more moles of magnesium chloride mixed with the same amount of water, more heat was released by the reaction. The experimental value was -151. 45 kJ/mol, which compared to the theoretical value of -155 kJ/mol3, there is a 2. 3% error. Any error could have been from temperature influences from the surrounding environment, inaccuracy in the calorimeter constant or slight human error in measurement.
Perhaps the calorimeter obtained was not clean or sufficient enough to reflect the constant more accurately. Another solution could have been the use of a better lid to keep the outside environment outside of the reaction. Again, more trials would have given more accurate conclusions and results, with each trial having smaller and less of a change in mass of magnesium chloride. If magnesium chloride were used as a deicer, the concrete would be affected by its chemical properties. Except compared to calcium chloride and sodium chloride, magnesium chloride is less severe in damages on concrete.
Magnesium chloride also isn’t harmful to vegetation and is the least toxic to vegetation of the chloride deicers4. Magnesium is often found in many fertilizers and is less toxic to the soil compared to calcium chloride and sodium chloride when they leach into surrounding soils. Luckily, magnesium chloride is only a mild skin and eye irritant in large quantities and the cost of magnesium chloride is approximately $0. 10 per kilogram. Conclusion The best experimental results for magnesium chloride’s ability to lower freezing point of water was 2. 88, meaning that magnesium chloride splits into almost 3 ions as it dissociates in water.
Experiment 3 did prove that magnesium chloride in water lowered the freezing point up to an average of 4. 516 degrees Celsius. Experiment 2 results returned an average enthalpy change of -151. 45 kJ/mol and demonstrated that magnesium chloride and water reaction is indeed exothermic. Based on the results of the experiments, as a deicer, magnesium chloride is a great choice because of its tendency to lower the freezing point of water, having a large Van’t Hoff factor, and the heat of solution is a pretty substantial amount compared to other salts such as, potassium chloride, which shows an endothermic reaction.
Since it is exothermic, magnesium chloride can lower the temperature while mixing with water before the enthalpy change wears off versus an endothermic reaction. Comparatively, magnesium chloride on its own would make the best road salt. The best road salt would lower the freezing point of water, be exothermic, coast effective, and non-harmful to the surrounding environment. Considering all the roadside salts have reasonably similar effects on the surrounding wildlife and they all do lower the temperature of water to some degree, the main factors looked at will be their enthalpy and cost effectiveness.
NaCl would make the second best deicer, even though it is an endothermic reaction with a heat of solution being 3. 9 kJ/mol5, the cost of one kilogram is $0. 07. Calcium chloride is an exothermic reaction, which in an ideal world would put it near the top of the list, but the cost relative to the other road salts is more higher, beyond its worth. Endothermic potassium chloride would follow with its cost about $0. 25 per kilogram and being the most expensive and least effective.
References
- Profmaster. com. “Heat Capacity with Dissolved Solids. ” (accessed October 17, 2012). ttp://profmaster. blogspot. com/2009/02/heat-capacity-with-dissolved-solids. html
- Wikipedia. “Solubility Table. ” (accessed October 18, 2012). http://en. wikipedia. org/wiki/Solubility_table
- Student Room. (accessed October 18, 2012). http://www. thestudentroom. co. uk/showthread. php? t=847759
- Ask The Builder. “Deicing Facts. ” (accessed October 18, 2012). http://www. askthebuilder. com/deicing-facts/
- WorldWideWolfe. “Enthalpies of Solution, Fusion, and Vaporization. ” (accessed October 18, 2012). http://www. mindspring. com/~drwolfe/WWWolfe_dat_enthalpies. htm
