All reactions involve some change in energy. During a physical change in matter, such as the evaporation of liquid water to water vapour, the energy of the water molecules increases. However, the change in energy is much smaller than in chemical reactions.

When a chemical reaction occurs, some bonds will break, while new bonds may form. Energy changes in chemical reactions result from the breaking and forming of bonds. For bonds to break, energy must be absorbed. When new bonds form, energy will be released because the new product has a lower energy than the `in between' stage of the reaction when the bonds in the reactants have just been broken.

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. This type of reaction is known as an exothermic reaction. 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. Most decomposition reactions are endothermic and heating is needed for the reaction to occur. Most synthesis reactions are exothermic, meaning that energy is given off in the form of heat or light.

More simply, we can describe the energy changes that take place during a chemical reaction as:

Total energy absorbed to break bonds - Total energy released when new bonds form

So, for example, in the reaction...

2Mg++O₂ →→ 2MgO

Energy is needed to break the O-O bonds in the oxygen molecule so that new Mg-O bonds can be formed, and energy is released when the product (MgO) forms.

Despite all the energy changes that seem to take place during reactions, it is important to remember that energy cannot be created or destroyed. Energy that enters a system will have come from the surrounding environment and energy that leaves a system will again become part of that environment. This is known as the conservation of energy principle.

Definition 1: Conservation of energy principle

Energy cannot be created or destroyed. It can only be changed from one form to another.

Chemical reactions may produce some very visible and often violent changes. An explosion, for example, is a sudden increase in volume and release of energy when high temperatures are generated and gases are released. For example, NH₄NO₃

The total mass of all the substances taking part in a chemical reaction is conserved during a chemical reaction. This is known as the law of conservation of mass. The total number of atoms of each element also remains the same during a reaction, although these may be arranged differently in the products.

We will use two of our earlier examples of chemical reactions to demonstrate this:

1. The decomposition of hydrogen peroxide into water and oxygen

2H₂O₂ →→ 2H₂O ++ O₂

Figure 1

Left hand side of the equation

Total atomic mass = (4 ×× 1) + (4 ×× 16) = 68 u

Number of atoms of each element = (4 ×× H) + (4 ×× O)

Right hand side of the equation

Total atomic mass = (4 ×× 1) + (2 ×× 16) + (2 ×× 16) = 68 u

Number of atoms of each element = (4 ×× H) + (4 ×× O)

Both the atomic mass and the number of atoms of each element are conserved in the reaction.

2. The synthesis of magnesium and oxygen to form magnesium oxide 2Mg++O₂→→2MgO

Figure 2

Left hand side of the equation

Total atomic mass = (2 ×× 24,3) + (2 ×× 16) = 80,6 u

Number of atoms of each element = (2 ×× Mg) + (2 ×× O)

Right hand side of the equation

Total atomic mass = (2 ×× 24,3) + (2 ×× 16) = 80,6 u

Number of atoms of each element = (2 ×× Mg) + (2 ×× O)

Both the atomic mass and the number of atoms of each element are conserved in the reaction.

In any given chemical compound, the elements always combine in the same proportion with each other. This is the law of constant proportion.

The law of constant composition says that, in any particular chemical compound, all samples of that compound will be made up of the same elements in the same proportion or ratio. For example, any water molecule is always made up of two hydrogen atoms and one oxygen atom in a 2:1 ratio. If we look at the relative masses of oxygen and hydrogen in a water molecule, we see that 94% of the mass of a water molecule is accounted for by oxygen and the remaining 6% is the mass of hydrogen. This mass proportion will be the same for any water molecule.

This does not mean that hydrogen and oxygen always combine in a 2:1 ratio to form H₂O. Multiple proportions are possible. For example, hydrogen and oxygen may combine in different proportions to form H₂O₂ rather than H₂O. In H₂O₂, the H:O ratio is 1:1 and the mass ratio of hydrogen to oxygen is 1:16. This will be the same for any molecule of hydrogen peroxide.

In a chemical reaction between gases, the relative volumes of the gases in the reaction are present in a ratio of small whole numbers if all the gases are at the same temperature and pressure. This relationship is also known as Gay-Lussac's Law.

For example, in the reaction between hydrogen and oxygen to produce water, two volumes of H₂ react with 1 volume of O₂ to produce 2 volumes of H₂O.

2H₂ + O₂→→2H₂O

In the reaction to produce ammonia, one volume of nitrogen gas reacts with three volumes of hydrogen gas to produce two volumes of ammonia gas.

N₂ ++3H₂ →→ 2NH₃

This relationship will also be true for all other chemical reactions.

The following video provides a summary of the concepts covered in this chapter.