All chemical reactions involve energy changes. In some reactions, we are able to observe these energy changes as either an increase or a decrease in the overall energy of the system.

When a chemical reaction occurs, bonds in the reactants break, while new bonds form in the product. The following example may help to explain this. Hydrogen reacts with oxygen to form water, according to the following equation:

In this reaction, the bond between the two hydrogen atoms in the H22 molecule will break, as will the bond between the oxygen atoms in the O22 molecule. New bonds will form between the two hydrogen atoms and the single oxygen atom in the water molecule that is formed as the product.

For bonds to break, energy must be absorbed. When new bonds form, energy is released. The energy that is needed to break a bond is called the bond energy or bond dissociation energy. Bond energies are measured in units of kJ.mol-1-1.

Definition 1: Bond energy

Bond energy is a measure of bond strength in a chemical bond. It is the amount of energy (in kJ.mol-1-1) that is needed to break the chemical bond between two atoms.

In some reactions, the energy that must be absorbed to break the bonds in the reactants, is less than the total energy that is released when new bonds are formed. This means that in the overall reaction, energy is released as either heat or light. This type of reaction is called an exothermic reaction. Another way of describing an exothermic reaction is that it is one in which the energy of the product is less than the energy of the reactants, because energy has been released during the reaction. We can represent this using the following general formula:

Reactants → Product + Energy Reactants → Product + Energy

Definition 2: Exothermic reaction

An exothermic reaction is one that releases energy in the form of heat or light.

In other reactions,the energy that must be absorbed to break the bonds in the reactants, is more than the total energy that is released when new bonds are formed. This means that in the overall reaction, energy must be absorbed from the surroundings. This type of reaction is known as an endothermic reaction. Another way of describing an endothermic reaction is that it is one in which the energy of the product is greater than the energy of the reactants, because energy has been absorbed during the reaction. This can be represented by the following formula:

Reactants + Energy → Product Reactants + Energy → Product

Definition 3: Endothermic reaction

An endothermic reaction is one that absorbs energy in the form of heat or light.

The difference in energy (E) between the reactants and the products is known as the heat of the reaction. It is also sometimes referred to as the enthalpy change of the system.

The heat of the reaction is represented by the symbol ΔHΔH, where:

Some of the information relating to exothermic and endothermic reactions is summarised in Table 2.

Definition 4: Enthalpy

Enthalpy is the heat content of a chemical system for a given pressure, and is given the symbol 'H'.

There are many examples of endothermic and exothermic reactions that occur around us all the time. The following are just a few examples.

Demonstration : Spontaneous and non-spontaneous reactions

Apparatus and materials:

A length of magnesium ribbon, thick copper wire and a bunsen burner.

Figure 1

Method:

1. Scrape the length of magnesium ribbon and copper wire clean.
2. Heat each piece of metal over the bunsen burner, in a non-luminous flame. Do Not look directly at the flame. Observe whether any chemical reaction takes place.
3. Remove the metals from the flame and observe whether the reaction stops. If the reaction stops, return the metal to the bunsen flame and continue to heat it.

Results:

• Did any reaction take place before the metals were heated?
• Did either of the reactions continue after they were removed from the flame?
• Write a balanced equation for each of the chemical reactions that takes place.

In the demonstration above, the reaction between magnesium and oxygen, and the reaction between copper and oxygen are both non-spontaneous. Before the metals were held over the bunsen burner, no reaction was observed. They need energy to initiate the reaction. After the reaction has started, it may then carry on spontaneously. This is what happened when the magnesium reacted with oxygen. Even after the magnesium was removed from the flame, the reaction continued. Other reactions will not carry on unless there is a constant addition of energy. This was the case when copper reacted with oxygen. As soon as the copper was removed from the flame, the reaction stopped.

Now try carefuly adding a solution of dilute sulfuric acid to a solution of sodium hydroxide. What do you observe? This is an example of a spontaneous reaction because the reaction takes place without any energy being added.

Definition 5: Spontaneous reaction

A spontaneous reaction is a physical or chemical change that occurs without the addition of energy.

From the demonstrations of spontaneous and non-spontaneous reactions, it should be clear that most reactions will not take place until the system has some minimum amount of energy added to it. This energy is called the activation energy. Activation energy is the 'threshold energy' or the energy that must be overcome in order for a chemical reaction to occur.

Definition 6: Activation energy

Activation energy or 'threshold energy' is the energy that must be overcome in order for a chemical reaction to occur.

It is possible to draw an energy diagram to show the energy changes that take place during a particular reaction. Let's consider an example:

Figure 2: The energy changes that take place during an exothermic reaction

The reaction between H2(g)H2(g) and F2(g)F2(g) (Figure 2) needs energy in order to proceed, and this is the activation energy. Once the reaction has started, an in-between, temporary state is reached where the two reactants combine to give H2F2H2F2. This state is sometimes called a transition state and the energy that is needed to reach this state is equal to the activation energy for the reaction. The compound that is formed in this transition state is called the activated complex. The transition state lasts for only a very short time, after which either the original bonds reform, or the bonds are broken and a new product forms. In this example, the final product is HF and it has a lower energy than the reactants. The reaction is exothermic and ΔΔH is negative.

Definition 7: Activated complex

The activated complex is a transitional structure in a chemical reaction that results from the effective collisions between reactant molecules, and which remains while old bonds break and new bonds form.

In endothermic reactions, the final products have a higher energy than the reactants. An energy diagram is shown below (Figure 3) for the endothermic reaction XY+Z→X+YZXY+Z→X+YZ. In this example, the activated complex has the formula XYZ. Notice that the activation energy for the endothermic reaction is much greater than for the exothermic reaction.

Figure 3: The energy changes that take place during an endothermic reaction