4 .4 Energy Changes Associated with Chemical Change — Endothermicity and Exothermicity

Warm Up

Complete the following table by placing a checkmark in the appropriate energy-change column for each of the classifications of chemical change listed. Reaction

Most Decomposition Reactions

Neutralization Reactions

Combustion Reactions

Sources of Energy

Fuels such as natural gas, wood, coal, and gasoline provide us with energy. Did you ever wonder where this energy comes from? Most of it is stored as chemical potential energy that can be converted into heat. The amount of potential energy available depends on the position of the subatomic particles making up a chemical sample and the composition of the sample. While the absolute potential energy content of a system is not really important, a change in potential energy content often is. We use changes in chemical energy to other energy forms to heat our homes and move our vehicles. Because chemical energy is most commonly converted to heat, we use the symbol ∆H to symbolize a change in energy available as heat. The symbol is sometimes read as delta H or an enthalpy change. Enthalpy is defined as potential energy that may be evolved or absorbed as heat.

All sorts of processes, both physical and chemical, have an enthalpy change associated with them. While a general change in enthalpy is symbolized as ∆H, specific types of enthalpy changes may be symbolized by a subscripted explanation attached to this symbol. Some examples include:

What is it that causes the enthalpy stored in reactants to differ from that in the products of a physical or chemical change? What exactly is responsible for producing a ∆H value? The answer is the chemical bonds.

Quick Check

Examine each of the equations listed on the previous page. Consider the sign of the ∆H value and determine whether the reactants or products have more stored energy in each case. Put a checkmark under the side with more enthalpy.

Reaction Reactants Products

Dissolving potassium hydroxide

Combustion of propane

Melting ice

Replacement of iron by aluminum

Formation of calcium hydroxide

The Energy of Chemical Bonds

Suppose your work text has fallen off your desk to the floor. When you lift it back to the desktop, you give the text potential energy because of its position above the floor. The gravitational potential energy your text now has could easily be converted into mechanical, as well as sound energy, should it happen to fall again.

In a similar way, the electrons in the atoms of the molecules of any substance have potential energy. Think of the negative electrons as being pulled away from the positive nucleus of their atom. If it were not for their high velocity, they would certainly rush toward and smash directly into the nucleus of their atom, much like your work text falling to the floor. By virtue of the position of the negative electrons in an atom relative to the positive nucleus and the other nuclei nearby in a molecule, the electrons (and the protons in the nuclei for that matter) have potential energy. This is the chemical potential energy called bond energy .

The bond energy just mentioned is the energy of intramolecular bonds. These are bonds that are formed between atoms within a molecule. Weaker bonds exist between molecules in a sample of solid, liquid, and even gaseous matter. These weak bonds hold the molecules of a solid or liquid together. These weak interactions between molecules are called intermolecular forces . The details of intermolecular forces relate to the polarity or lack of polarity of a molecule. The difference between the potential bond energy of reactants and products before and after a chemical or physical change is known as the enthalpy change or ∆H value.

Bond breaking requires energy input while bond forming results in energy release.

It’s fairly intuitive to most of us that breaking something requires energy. For a karate master to break a board, he must apply energy. Similarly breaking chemical bonds requires energy. This energy is used to overcome the electrostatic forces that bind atoms or ions together. It follows from this that bond formation must result in a more stable situation in which energy is released.

If more energy is absorbed as bonds break than is released as bonds form during a chemical change, there is a net absorption of energy. As energy is required to break bonds there will be a gain in enthalpy. Reactions such as these have a positive ∆H value and are called endothermic reactions. Such a reaction may be shown graphically in an energy profile, sometimes called a potential energy diagram (Figure 4.4.1).

Quick Check

Potential Energy vs. Reaction Proceeding

ΔH > 0

Reaction Proceeding

The energy absorbed to break bonds is greater than the energy released during bond formation in an endothermic reaction so they have positive ∆H values.

If, on the other hand, more energy is released as bonds are formed than is absorbed as they are broken during a chemical change, there is a net release of energy and so a decrease in enthalpy content occurs. Reactions such as these will have a negative ∆H value and may be called exothermic. Such a reaction may be represented graphically as shown in Figure 4.4.2.

Potential Energy vs. Reaction Proceeding

reactants

∆

Reaction Proceeding

More energy is released during bond formation than is absorbed during bond breaking in an exothermic reaction so they have negative ∆H values.

Complete the right column in the following table by labelling each of the introductory reactions in this section as endothermic or exothermic.

Reaction

Dissolving potassium hydroxide

Combustion of propane

Melting ice

Replacement of iron by aluminum

Formation of calcium hydroxide

Endothermic or Exothermic