0.9 Energetics of chemical reactions

The foundation

We begin our study of the energetics of chemical reactions with our understanding of mass relationships,determined by the stoichiometry of balanced reactions and the relative atomic masses of the elements. We will assume a conceptualunderstanding of energy based on the physics of mechanics, and in particular, we will assume the law of conservation of energy. Indeveloping a molecular understanding of the reaction energetics, we will further assume our understanding of chemical bonding viavalence shell electron pair sharing and molecular orbital theory.

Goals

The heat released or consumed in a chemical reaction is typically amongst the most easily observed and mostreadily appreciated consequences of the reaction. Many chemical reactions are performed routinely specifically for the purpose ofutilizing the heat released by the reaction.

We are interested here in an understanding of the energetics of chemical reactions. Specifically, we wish to knowwhat factors determine whether heat is absorbed or released during a chemical reaction. With that knowledge, we seek to quantify andpredict the amount of heat anticipated in a chemical reaction. We expect to find that the quantity of heat absorbed or releasedduring a reaction is related to the bonding of the molecules involved in the reaction.

Prior to answering these questions, we must first answer a few questions regarding the nature of heat. Despiteour common familiarity with heat (particularly in Houston), the concept of heat is somewhat elusive to define. We recognize heat as"whatever it is that makes things hot," but this definition is too imprecise to permit measurement or any other conceptual progress.Exactly how do we define and measure heat?

Observation 1: measurement of heat by temperature

We can define in a variety of ways a temperature scale which permits quantitative measurement of "howhot" an object is. Such scales are typically based on the expansion and contraction of materials, particularly of liquid mercury, or onvariation of resistance in wires or thermocouples. Using such scales, we can easily show that heating an object causes itstemperature to rise.

It is important, however, to distinguish between heat and temperature. These two concepts are not one andthe same. To illustrate the difference, we begin by measuring the temperature rise produced by a given amount of heat, focusing onthe temperature rise in 1000g of water produced by burning 1.0g of methane gas. We discover by performing this experiment repeatedlythat the temperature of this quantity of water always rises by exactly 13.3°C. Therefore, the same quantity of heat mustalways be produced by reaction of this quantity of methane.

If we burn 1.0g of methane to heat 500g of water instead, we observe a temperature rise of 26.6°C. Ifwe burn 1.0g of methane to heat 1000g of iron, we observe a temperature rise of 123°C. Therefore, the temperature riseobserved is a function of the quantity of material heated as well as the nature of the material heated. Consequently, 13.3°Cis not an appropriate measure of this quantity of heat, since we cannot say that the burning of 1.0g of methane "produces13.3°C of heat." Such a statement is clearly revealed to be nonsense, so the concepts of temperature and heat must be keptdistinct.