Qualitative and Quantitative Determination of ∆H,

the Heat of Reaction

Experiment # 7: Calorimetry
Affiliation: PCC
Chemistry 222

ABSTRACT.

In this experiment, the temperature changes that occurred during two processes were determined. The first process was the dissolution of a salt in water. The second process was the neutralization of an acid with a base. The experimental data gathered allowed the calculation of the sign of the heat of solution and the heat of neutralization. Many salts dissolve in water to form aqueous solutions of the salts. This process is accompanied by a change in energy in the form of heat. Similarly when an acid and base react in a neutralization reaction, the water formed is accompanied by a change in energy also in the form of heat. These reactions were performed in the lab, open to the atmosphere (i.e. at constant pressure), allowing calculation of this energy change. The process by which these calculations were performed and the measurement of this energy completed is called calorimetry. The values obtained during these experiments were compared to theoretical values and it was noted that the qualitative values were as expected while the quantitative values they were significantly larger than expected as they were outside the acceptable 5% margin of error.

Introduction

Chemical systems contain a multitude of particles, each containing potential and/or kinetic energy. Internal energy of a system is defined as the sum of all these energies (Silderberg, 2009). When reactions occur in chemical systems, the chemical changes are accompanied by changes in energy, usually in the form of heat. The enthalpy change, ∆H, also called the heat of reaction, occurs at constant pressure and is the most important part energetics, the study of energy under transformation. For reactions where heat is evolved, the reactions are described as exothermic with ∆H0.

Heats of Solution

When a solute and solvent are combined, the two contain particles which attract each other. In order for one substance to dissolve another, three things must occur: (1) the solute particles must separate from each other, (2) some solvent particles must separate to make room for the solute particles, and (3) solute and solvent particles must mix together (Silderberg, 2009). When this occurs some energy must be absorbed and some energy must be released. Ionic compounds are made up of crystal lattice structures and the attraction between the ions in this structure is the crystal lattice energy. The energy required to break this crystal lattice is known as ΔHLattice. Since energy is added to the system in order to break the lattice, this is an endothermic process and ΔHLattice is always positive (ΔHLattice>0). When the ions become surrounded by water and are solvated intermolecular bonds between the ions and water form. This process releases energy known as the energy of hydration, ΔHHydration. Since energy is released from the system, the process is exothermic and ΔHHydration is always negative (ΔHHydration ΔHHydration then ΔHSolution is positive and the reaction is endothermic and if ΔHLattice LiCl(aq) + H2O(l)
More energy was released during the hydration process and formation of aqueous salt than was required to break the bonds in the lattice structure. As a result, ΔHHydration > ΔHLattice and this process was exothermic. In contrast, NH4Cl(s) dissolved to give NH4Cl(aq) which caused the test tube to feel cool to touch as heat was absorbed from the surroundings in the following reaction:
NH4Cl(s) + H2O(l) ----> NH4Cl(aq) + H2O(l)
This process required more energy to break the bonds in the lattice structure due to the presence of H-bonds and dipole forces. As a result, ΔHLattice > ΔHHydration and the process was endothermic as heat was drawn from the surroundings in order to complete this reaction. These results were as expected due to the molecule sizes and bond