# Thermochemistry formal lab

Topics: PH, Enthalpy, Hydrochloric acid Pages: 6 (2337 words) Published: April 26, 2014
﻿Thermochemistry Laboratory Report
Abstract
The purposes of these three experiments are to determine the heat capacity of a calorimeter and with that data, confirm Hess’s Law and observe enthalpy changes within reactions. By measuring the change in temperature that occurs with the interaction of two different reactants, we were able to determine both the calorimeter constant and the change in enthalpy of a given reaction. The results were rather mixed, as some numbers more closely resembled the theoretical values than others did.

Introduction
The first experiment is devoted to finding the calorimeter constant for a polystyrene cup. Whenever a reaction takes place inside a calorimeter, some heat is lost to the calorimeter and its surroundings. In order to achieve maximum accuracy, we must know exactly how much heat will be lost, so that the results of the next two experiments will be as correct as possible. The equation used to determine it is a simple manipulation of the overall heat of the reaction equation, which is: Overall Heat = - [(Sp.Ht. hotwater * Mass of water * Change in temperature) + (Sp.Ht. coolwater * Mass of water * Change in temperature) + (Cp calorimeter * Change in temperature)] Since an error is bound to happen during the experimental process, three calculations were done to find an average. This experiment is vital to the success of the following two thermochemistry experiments.

The second experiment, entitled Hess’s Law, is a simple confirmation of said law. To do so, we take three reactions, where one of them is the same as the other two, and measure the heats of reaction for each of them. Hess’s Law states that the heat of reaction of the one reaction should equal to the sum of the heats of reaction for the other two. The three reactions used in this experiment are: (1) NaOH(s)  Na+(aq) + OH-(aq)

(2) NaOH(s) + H+(aq) + Cl-(aq)  H2O(l) + Na+(aq) + Cl-(aq) (3) Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq)  H2O(l) + Na+(aq) + Cl-(aq) In order to find the heat released by each reaction, we used a variant of the overall heat of a reaction equation, which was q = - [Sp.Ht. * m * Change in temp.]. In addition to finding the change in enthalpy, change in entropy was also calculated using theoretical values in given reference tables. Finally, the overall free energy released was calculated using the equation: Change in free energy = Change in enthalpy – (Temperature * Change in entropy). All of this is then used to verify Hess’s Law by calculating the percent error involved in the experiment.

The third experiment, called Thermochemistry: Acid + Base, combines the concepts of the previous two experiments. The main concept is to observe the change in enthalpy that results from the various reactions between strong and weak acids and bases. There were four reactions used in this experiment, and they are: (1) HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)

(2) HCl(aq) +NH3(aq)  NH4Cl(aq)
(3) HC2H3O2(aq) + NaOH(aq)  NaC2H3O2(aq) + H2O(l)
(4) HC2H3O2(aq) + NH3(aq)  NH4C2H3O2(aq)
By monitoring the change in temperature that results from the reaction of an acid and a base, it is possible to calculate the overall energy for each reaction, also known as ∆H rxn/mole of limiting reactant. This experimental value can be compared with the theoretical to determine how accurate the experiment was. The lower the percent error, the more accurate we were at calculating the energy involved in each reaction.

Experimental
In order to do any calculation for energy, we first had to find the calorimeter constant. In order to do that, we first took and weighed a polystyrene cup (our calorimeter) and added approximately 100 g of warm water to it. The actual measurements are recorded in Table 1-1. The mass of the cup with the water in it were recorded to find the exact mass of the water added. Next, a cylinder was weighed, like the cup, and about 48 mL of cool water was added. The total was weighed and recorded in the same table....