0.12 Reaction energy and bond energy

Introduction

In the previous study, we developed detailed means to observe and measure the energy changes in chemical reactions. This ability is valuable all on its own, since managing the flow of energy from one form to another is a vital economic activity. But our work is not half done. In Chemistry, we seek not just to observe and measure, but also to model and to understand conceptually. We can make this point clearly by thinking about the following. Some chemical reactions produce energy, even in spectacular amounts. The detonation of a single gram of trinitrotoluene (TNT) produces about 4.2 kJ of energy. The reaction of a single gram of sodium metal (Na) with water produces about 8 kJ of energy. On the other hand, some reactions absorb energy, often evidenced by a significant cooling of the products or the surroundings of the reaction. For example, the hydration of ten grams of ammonium nitrate (NH ₄ NO ₃ ) in an instant cold pack absorbs about 3.2 kJ of energy, causing it to be cold enough to treat minor athletic injuries.

How can we account for these great variations in the energies of reactions? Where does the energy come from in an exothermic reaction, and where does it in an endothermic reaction? Could we find a way to predict whether a reaction will be exothermic or endothermic? Answering these questions requires us to develop a model for energy transfer during chemical reactions.

Foundation

We will build significantly on the results of the previous concept study. We know how to measure energy changes in reactions. A reaction which releases energy into the environment is called an exothermic reaction, and the heat transfer q<0. A reaction which absorbs energy from the environment is called an endothermic reaction, and the heat transfer q>0.

Hess’ Law, developed in the previous concept study, is an extremely important observation. Recall that Hess’ Law tells us that the energy of a reaction is equal to the sum of the energies of a set of reactions which add up to the overall reaction. Stated differently, the energy of a reaction does not depend on what “path” we follow in converting reactant to products, whether it be in a single reaction or a series of reactions. As long as we start with the same reactants and wind up with the same products, the energy of the reaction is the same.

Although this is not an observation or previous conclusion, we’ll add to our foundation a definition of a new quantity, called “enthalpy.” To understand the usefulness of this new quantity, let’s remember that, according to Hess’ Law, if we start with a set of reactants and carry out a series of reactions which recreate the reactants, then the total energy change summed over that series of reactions has to be exactly zero. Using the Law of Conservation of Energy, this makes sense. We would not expect to be able to change the energy of a substance or substances without changing the state of those substances. In fact, for this reason, chemists call the energy of a substance a “state function,” meaning that the energy depends only on what state the substance is in (gas, liquid, solid; temperature; pressure).