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Heats of reaction and Hess� Law Abstract

Heats of reaction and Hess� Law Abstract - follows...

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Abstract: Heats of reaction and Hess’ Law In a calorimetric experiment such as this, both the liberation and consumption of heat is measured and analyzed by the rise and fall of temperature of the system (defined as the reactants in the calorimeter). This is made possible by the fact that energy can neither be created, nor destroyed (First Law of Thermodynamics). In this experiment, the purpose was to calculate the heats and enthalpies of reaction for four acid-base pairs: (a) HCl-NaOH, (b) CH 3 COOH-NaOH, (c) HCl-NH 3 , and (d) CH 3 COOH-NH 3 with calorimetric analysis. The heat gained by the calorimeter per degree Celsius temperature change was 51.5 J/gºC. From this, the heats of reactions (a)-(d) were calculated along with the enthalpies of reactions (a)-(d), which are as
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Unformatted text preview: follows: ΔH rxn, (a) = -59.8 kJ/mol, ΔH rxn, (b) = -53.6 kJ/mol, ΔH rxn, (c) = -57.5 kJ/ mol, and ΔH rxn, (d) = -56.2 kJ/mol. Using these experimentally-derived enthalpy values and setting up a Hess’ Law problem, a value for the enthalpy of the neutralization for (d) was calculated to be -51.3 kJ/mol, which constituted an 8.7% error from the directly-calculated value. Considering this relatively small error value, it can be concluded that the experiment was conducted well and that the theory behind calorimetry, including the Laws of Thermodynamics, are accurate for this experiment....
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