Name: Victoria BlacklockDetermining the Molar Enthalpies of Calcium Chloride, Potassium Nitrate and Potassium Chloride Using Calorimetry October 18th, 2018Introduction:Calorimetry is the experimental process used to solve for the amount of heat absorbed or released within the system or surroundings during a chemical or physical change. Calorimetry is used to find the Molar enthalpy of a reactant or product. In this case the Molar enthalpy of solution is the total amount of thermal energy released or absorbed in the dissolving process at constant a pressure corresponding to the amount of moles in the compound. The purpose of this experiment is to measure the amount of thermal energy needed to dissolve calcium chloride, potassium nitrate and potassium chloride in water.
Materials and Methods: Refer to lab sheet and p.347 of the Nelson Chemistry 12Observations:Table 1: Calorimeter Data of CaCl2, KNO3 and KClSalt Trial Mass (g) Temperature Change (C)CaCl2 Trial 1 4.641 8.9Trial 2 4.931 10.4Trial 3 3.
786 6.2KNO3 Trial 1 5.152 -7.9Trial 2 5.000 -4.1Trial 3 4.161 -4.
4KClTrial 1 4.409 -6.0Trial 2 5.161 -6.3Trial 3 4.
338 -4.0Table 2: Average Mass and Temperature Change of CaCl2, KNO3 and KClSalt Mass(g) Temperature Change (C)CaCl2 4.453 8.5KNO3 4.771 -5.
5KCl4.636 -5.4Data Manipulations: CaCl2:HSol= -( 4.18KJKg C x 0.0500Kg1 x 8.5C1 x 110.
98gmol x 14.453g) = -44 KJ/mol% Error= (-81.30-(-44)-81.30)x 100 = 46%KNO3:HSol= -( 4.18KJKg C x 0.0500Kg1 x -5.
5C1 x 101.11gmol x 14.771g) = 24 KJ/mol% Error= (34.88-2434.88)x 100 = 31%KCl:HSol= -( 4.18KJKg C x 0.0500Kg1 x -5.
4C1 x 74.55gmol x 14.636g) = 18 KJ/mol% Error= (18-17.2217.22)x 100 = 5.0%Results Summary:The accepted molar enthalpy of calcium chloride is -81.30KJmol , the experiment generated a molar enthalpy of -44KJmol . This signified that the calculations of calcium chloride were 60% less than expected, leading to a 46% error.
Potassium nitrate has an accepted molar enthalpy of 34.88KJmol , with calculations the experiment generated a molar enthalpy of 24KJmol. This signified that the calculations were 37% less than the accepted molar enthalpy, causing there to be a 31% error. Potassium chloride has an accepted molar enthalpy of 17.22KJmol, while the experimental molar enthalpy is 18KJmol. This signified that the calculations were 4.4% more than the accepted molar enthalpy, generating a 5.
0% error. Calcium chloride has the largest percent error and potassium chloride having the smallest. Discussion: Each reaction has a varying molar enthalpy, this is because each reaction needs a different amount of energy to break the bonds within the reactants (absorbed), or make new bonds within the products (released).
In the dissociation of calcium chloride, potassium nitrate and potassium chloride, the water molecules are attracted to and bond to the salts. Potassium nitrate solution and potassium chloride solution both have a positive molar enthalpy, meaning that both of these reactions are endothermic. Endothermic reactions are chemical, physical or nuclear reactions that occur when the amount of energy absorbed is greater than the amount released. This energy is used to break the bonds within the reactants. In this case, the amount of energy used for the water molecules to bond to the solute is less than the amount needed to separate it.
This indicates that there is a temperature decrease in the surroundings, since the reaction absorbs energy. Calcium chloride is the only solution in this experiment that has a negative molar enthalpy, this signifies that this reaction is exothermic. Exothermic reactions are chemical, physical, or nuclear reactions that occur when the amount of energy released is greater than the amount absorbed.
This energy is used to make the bonds. In this case, the greater amount of energy is released when the water molecules bond to the solute. This indicates that there is an increase of temperature in the surroundings, since the reaction releases energy. The accepted and the experimental molar enthalpies results are different based on the assumption that no heat was transferred between or lost to the surroundings.
If this assumption was true the experimental molar enthalpies would have been much more accurate than precise.
