Heats Of Reaction Chemistry 141 220 Ta Olivia Andrus 111419

Heats Of Reactionchmy 141 220ta Olivia Andrus111419introductionas T

Heats Of Reactionchmy 141 220ta Olivia Andrus111419introductionas T

Paper For Above instruction

This paper examines an experiment conducted to explore the heats of various chemical reactions using thermodynamic principles, notably Hess’s Law. The main goal was to measure the heat released during reactions of potassium hydroxide (KOH) with water and hydrochloric acid (HCl), and to verify whether the heats of combined reactions align with theoretical predictions based on Hess’s Law. The experiment employed thermistors to record temperature changes, and the data collected was analyzed to calculate the enthalpy changes and evaluate the validity of Hess’s Law in this context.

The experiment involved three key reactions. In the first, solid KOH was dissolved in water; in the second, solid KOH was reacted with HCl; and finally, the KOH solution from part one was subsequently reacted with HCl to observe if the total heat released matched the sum of the individual reactions. Thermistors suspended over styrofoam cups containing the reacting solutions recorded temperature changes, which were then graphed and analyzed. These temperature data points provided the basis to compute the heat evolved during each reaction.

The initial setup involved measuring precise quantities of KOH pellets, water, and HCl solutions, ensuring accurate calculations of moles involved. The temperature data captured before and after adding reactants allowed for calculation of temperature change (ΔT). Using specific heat capacities and the mass of the solutions, the heat released (Q) was calculated for each reaction with the formula Q = mcΔT, where 'm' is the mass, 'c' is the specific heat capacity of water (4.184 J/g°C), and ΔT is the temperature change. The number of moles of reactants was determined from the measured masses and molar masses, facilitating the determination of molar heats of reaction (enthalpies).

The experimental results showed that each reaction was exothermic, evidenced by increases in temperature. The first reaction, KOH in water, produced the highest heat release, followed by the second, involving KOH and HCl, with the third – the mixed reaction – yielding the least. Graphs plotted from the temperature data demonstrated characteristic sudden increases at the moment of reactant addition, confirming the exothermic nature of these reactions. However, the calculations of enthalpies revealed significant disparities compared to theoretical values, with percent errors exceeding 1000% in some cases. These discrepancies suggest experimental errors or incomplete reactions, such as insufficient dissolution of KOH or heat loss to the surroundings.

Applying Hess’s Law, it was expected that the heat of the third reaction (the combination of parts one and two) would equal the sum of the heats recorded in parts one and two separately. While the theoretical calculations support this principle, the experimental data did not align, showing a high percent error (~1438%). Several factors could have contributed to this failure, including heat dissipation, measurement inaccuracies, or reaction inefficiencies. The loss of heat to the environment, incomplete dissolution of reactants, or contaminated solutions could have affected the thermistor readings, leading to inaccurate calculations.

In conclusion, the experiment successfully demonstrated the exothermic nature of the reactions and the utility of thermistors in measuring temperature changes. It also aimed to verify Hess’s Law; however, due to significant experimental inaccuracies and errors, the data did not support the law’s predictions in this case. To improve accuracy, future experiments should consider better insulation, more precise measurement techniques, and ensuring complete reaction of all reagents. Despite the limitations, the experiment highlighted the principles of thermodynamics and the importance of meticulous experimental procedures to validate fundamental chemical laws.

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