Determining enthalpies of reaction by calorimetry

Apar G.

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Determining enthalpies of reaction by calorimetry

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Abstract

The purpose of this lab is to determine the specific heat of the calorimeter, the enthalpy changes of 4 different reactions, namely Heat of Solution of NaOH(s), Heat of Solution of KNO3(s), the Heat of Neutralization of HCl(aq) and NaOH(aq), and the Heat of Reaction of NaOH(s) and HCl(aq), and to verify Hess's Law. This will be determined by using calorimetry, which measures the heats of reaction. This relies on the First Law of Thermodynamics, which applies the theorem of conservation of energy to heat processes. In this lab, the calorimeter consists of 2 nested Styrofoam cups, which insulate it from the surroundings and allow it to be an isolated system.1 The values obtained will be compared to literature values and will also help to determine the heat capacity of the calorimeter.

Introduction
The equation
Sample calculation
Discussion
Conclusion
References

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[...] In order to verify Hess's Law, the enthalpies for the individual steps for the reaction are summed and compared with the change in enthalpy between the final and initial states for the overall reaction. Therefore as reaction one and three are steps in the overall reaction (denoted by reaction the ΔH for reaction and three are summed and compared with the overall ΔH for reaction four. The sum of ΔH for reaction one and three are - 45.34 kJ+ - 59.18 kJ= - 104.52 kJ. [...]

[...] The chemical equation for this reaction is as follows: KNO3(s) + H2O + NO3- Total Mass of Solution Heat Change in Enthalpy Change (ΔH [Kelvin] ΔT+CcalΔT) Table Calculations for qreaction/H2O for KNO3 qKNO3 = - q H2O = 2.2321 kJ/mol ΔH = q KNO3/n KNO3; n = ( 4.26 g KNO3) /(101g KNO3/mol) = 0.0422 mol ΔH = 2.2321 kJ / 0.0422 mol = 52.89 KJ/mol. Trial ΔH (KJ/mol) 1 52.89 2 45.26 Table Heats of Reaction for Reaction 2 Average ΔH for Reaction 49.08 KJ/mol Heat of Neutralization (Reaction Aqueous HCl and NaOH are mixed and the temperature change was plotted in a temperature vs. [...]

[...] This is shown as the temperature recordings for the NaOH reaction show a slight increase over time as the temperature re-equilibrates with the temperature of the surroundings. Reaction Three and Four's relatively low percent error could also be accounted for this minor heat loss due to the inherent imperfectness of the calorimeter insulating system. However, Reaction Two's large percent error ( 106.3 could be due to the fact that because this particular reaction is an endothermic reaction, the previous exothermic reactions could have caused heat to be released from the Styrofoam cups into the system, artificially adding energy and increasing the value of the magnitude of the change in energy. [...]

[...] In addition, the total heat of the reaction must be divided by the number of moles substance reacting in order to calculate the enthalpy change of the reaction (ΔH =qreaction/number of moles substance reacting). If the system is isolated, as present in this experiment, then the water will absorb the heat of the reaction = qreaction). The equation can be utilized in order to calculate the heat of reaction. However, for the best accuracy, we must account for the fact that the colorimeter itself may absorb some heat. [...]

[...] Results Calculation of Heat Capacity of Coffee-Cup Calorimeter Tcold water= 23.2 Density of water = 1 g/mL Thot water= 84.5 If a known mass of hot water is added to a known mass of cold water, then the following relation must hold true: Heat lost by hot water = Heat gained by cold water + heat gained by calorimeter In order to determine Ccal, hot water at approximately 90°C was added to cold water at approximately 25°C. The results are plotted in Graph 1. [...]

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