Exothermic and endothermic reactions

All chemical reactions involve an energy change. We categorise reactions by the direction of this energy change. If energy is released by a reaction, the reaction is exothermic. If energy is absorbed by a reaction, the reaction is endothermic.

When petrol burns via a combustion reaction, the energy released can be used to power a car. Combustion reactions release energy to the environment and so are exothermic reactions.

When a portable ice pack is activated for a sports injury, the reaction inside absorbs energy from the region around the injury where it is placed. By absorbing heat energy from the surroundings, the pack feels cold to the touch. Reactions that absorb energy from the surroundings are endothermic reactions.

The amount of energy change in a reaction can be determined. For combustion of fuel s, determining this energy change can allow the energy density of fuels to be compared.

Use this page to revise the following concepts within exothermic and endothermic reactions:

• Enthalpy
• Combustion
• Writing combustion equations
• Thermochemical equations

Enthalpy

The unique arrangement of particles and charges in each substance leads to it having a unique amount of chemical energy. Chemical energy is also called enthalpy, denoted by the symbol ΔH. Every chemical reaction will have a change in enthalpy, ∆H, as reactions involve changes in the arrangement of atoms.

In general, during a chemical reaction, the change in enthalpy can be represented as:

Exothermic reactions

There is always a net release of energy in exothermic reactions. However, before energy can be released, energy must first be absorbed by the reaction to break the bonds between the reactants. The energy needed to break the bonds is known as the activation energy, E_a.

The energy changes in a reaction can be represented in an energy profile diagram .

In an exothermic reaction:

• The energy released when the bonds in the product form is greater than the energy required to break the bonds in the reactants.
• The combined enthalpy of the products is less than the combined enthalpy of the reactants.
• ∆H will be negative as H_P < H_R.
• Energy is released to the surroundings.

Endothermic reactions

In an endothermic reaction:

• The energy absorbed by the reactants is greater than the energy released when the products form.
• The combined enthalpy of the products is greater than the combined enthalpy of the reactants.
• ∆H will be positive as H_P > H_R.
• Energy is absorbed from the surroundings.

Combustion

Combustion reactions are used to obtain energy from many fuels. Combustion reactions are exothermic reactions in which the fuel combines with oxygen, releasing energy.

Complete combustion will usually occur when the supply of oxygen gas is abundant. Some examples are:

If the amount of oxygen is limited, incomplete combustion occurs, forming carbon or carbon monoxide, CO. The equation below shows incomplete combustion of propane to form carbon monoxide and water.

C₃H₈(g) + 3.5O₂(g) → 3CO(g) + 4H₂O(l)

Writing combustion equations

The simplest way to balance a combustion equation is to balance the atoms in the order: carbon, hydrogen, then oxygen.

Example: Writing a combustion equation for butane

Thermochemical equations

Thermochemical equations can be used to compare fuels. Thermochemical equations include a value for the energy change that occurs in a reaction. This value includes a positive or negative sign to indicate if the reaction is exothermic or endothermic . The thermochemical equation for the combustion of methane is:

CH₄(g)   +  2O₂(g)   →   CO₂(g)   +  2H₂O(l)            ∆H   =- 890 kJ

The above equation indicates:

• The reactant and product molecules and their states.
• The reaction is exothermic.
• The reaction between 1 mole of methane and 2 moles of oxygen releases 890 kJ of energy.
• Methane is a useful fuel as 890 kJ is a large value.

Variations of a thermochemical equation and enthalpy change

The change in enthalpy of a reaction is linked to the stoichiometry of a reaction, and the states of reactants and products. If these are varied, then ΔH will change accordingly.

Energy density

Consider the combustion reactions:

CH₄(g)   +  2O₂(g)   →   CO₂(g)   +  2H₂O(g)         ∆H = -809 kJ mol^-1

C₃H₈(g)   +  3.5O₂(g)   →   3CO(g)   +  4H₂O(l)     ∆H = -2220 kJ mol^-1

The \(\Delta H\) for the above equations might suggest propane is a far better fuel than methane. However, it is comparing 16 g of methane (1 mole) with 44 g of propane (1 mole). When compared per gram, methane has a higher energy density.

∆H g^-1 for methane \(\displaystyle = \frac{890\text{ kJmol}^{-1}}{16\text{ gmol}^{-1}} = 55.6\text{ kJg}^{-1}\) or \(55.6\text{ MJkg}^{-1}\)

∆H g^-1 for propane \(\displaystyle = \frac{2220\text{ kJmol}^{-1}}{44\text{ gmol}^{-1}} = 50.5\text{ kJg}^{-1}\) or \(50.5\text{ MJkg}^{-1}\)

Energy density can be displayed in several ways. For methane:

• molar enthalpy of combustion ∆H = - 890 kJ mol^-1 (only enthalpy values have a negative sign)
• molar heat of combustion ∆H = 890 kJ mol^-1
• heat of combustion = 55.6 kJ g^-1

Note that the negative sign is only applicable to the enthalpy. The molar enthalpy of combustion and the heat of combustion of some fuels are compared in the table below.

Examples of energy calculations

Energy released by a fuel can be calculated using the formula

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