Lab Discussion -Experiment 9: Thermochemistry
In this experiment, we have the opportunity to observe the role thermodynamics plays in chemical reactions and all of the mechanics that pertain with how it works. The objectives of this experiment were accomplished through the use of calorimetry, the acknowledgment of endothermic and exothermic rules, and through the use of calorimetric equations and knowledge of thermochemistry.
In order to observe how temperature, heat capacity, and heat interact with each other through the process of calorimetry, the calorimeter must first be calibrated. In Part one of this experiment, the calorimeter is initially set up and the heat capacity of the calorimeter is found by noting the temperature difference between the added cold and hot water. Afterwards, you can utilize the Ccal value along with the temperature changes observed by the calorimeter to find the specific heat capacity of an unknown metal. This is done through adding room temperature water to the calorimeter and then adding the unknown metal to the cup. The specific heat capacity of the metal can then be found through this equation:
(mc)(cs)(ΔTc) + Ccal(ΔTc) = -(m)(cmetal)(ΔTH)
A similar process will also take place in order to find the heat of solution and the heat of neutralization. In order to obtain the heat of solution in this experiment, which is the enthalpy change associated with the dissolution of a substance in a solvent at constant pressure, room temperature water is added to the calorimeter and then NH4No3 is added. The desired molar heat of solution may be found after taking the temperature values, the average Ccal value, and the moles of NH4NO3 and properly plugging the numbers into the following equations:
q = – (CcalΔT + msolutionCsΔT)
In order to find the molar heat of neutralization in this experiment, which is the change in enthalpy that occurs when one equivalent of an acid and one equivalent of a base undergo a neutralization reaction to form water and a salt, you must add NaOH to the calorimeter and later add HCl. In order to calculate the molar heat of neutralization, take the proper temperature values, the average Ccal value, and the moles of H2O produced and properly plug the numbers into the following equations:
q = – (CcalΔT + msolutionCsΔT)
The main function associated with the calorimeter is that of measuring differences in heat. Calorimeters must be calibrated due to general concept of how heat is transferred from hot to cold. Essentially the amount of heat that leaves a reaction is equivalent to the amount of heat that is absorbed by both the calorimeter and solution. Calorimeters have the ability to absorb and release heat, thus they must be calibrated before use. Ccal represents the calorimeter’s specific heat capacity and it is measured so that later on in the lab the heat capacity of an unknown metal may be found and so the heat of a solution and neutralization can also be obtained. It is also primarily measured so that you can know how much heat is flowing to either the calorimeter or the other substance. The Ccal value is measured twice predominantly because it allows you to ensure that the value is authentic. By having a value that is not accurate, you risk obtaining flawed results for the rest of the experimentation. The first Ccal value we obtained was 47.652J/k and the second Ccal value was 69.95J/K. Our results were reproducible. We did not acquire the same exact Ccal value, however both of the values that were obtained were within the proper range of 10-80J/K. The average of the procured values is 58.725J/K.
The unknown metal was found to be Zinc. The Cmetal value we obtained was 0.3765 J/Kg a value that was not close enough to the given Zinc specific heat capacity value. The percent error was 3.1 percent. A possible source of error could be that we waited too long to place the metal into the calorimeter which could lead to a notable change in the resulting specific heat capacity value. The atomic mass we obtained was 66.4 g/mol while the actual atomic mass is 64.27 g/mol. The percent error between these values is 3.31 percent. Another source of error for part two could be associated with an error in the weighing of the unknown metal. If too much metal or too little of it was weighed, that would cause a resulting difference in temperature, thus the calculated values would be slightly off as well.
In Part 3, the endothermic reaction, the value of the experimental change in enthalpy of the solution was found to be 34.086 kJ/mol. While the value of the calculated change in enthalpy of the solution found by using Hess’s law was 28.6 kJ/mol. The percent error between these values is 19.18 percent. The net ionic reaction for Part 3 is as follows:
NH4NO3(s) →NH4+ (aq) + NO3- (aq)
A possible source of error for part three could be associated with how we could have waited a bit too long to place the ammonium nitrate into the calorimeter which could have a significant impact regarding what the final temperature would be.
In Part four, the exothermic reaction, the value of the experimental change in enthalpy of the solution was found to be -60.38 kJ/mol, while the value of the calculated change in enthalpy of the solution was -55.83 kJ/mol. The percent error these values is 8.15 percent. A source of error in part four could be associated with how the temperatures of HCl and NaOH could have been more that 0.2 degrees Celsius off from each other which could have had a significant impact in the data that we obtained, thus our final temperature and final value for the experimental change in enthalpy would be off.