Bond Energies - Chemistry LibreTexts

Bond Breakage and Formation

When a chemical reaction occurs, the atoms in the reactants rearrange their chemical bonds to make products. The new arrangement of bonds does not have the same total energy as the bonds in the reactants. Therefore, when chemical reactions occur, there will always be an accompanying energy change.

In some reactions, the energy of the products is lower than the energy of the reactants. Thus, in the course of the reaction, the substances lose energy to the surrounding environment. Such reactions are exothermic and can be represented by an energy-level diagram in Figure 1 (left). In most cases, the energy is given off as heat (although a few reactions give off energy as light). In chemical reactions where the products have a higher energy than the reactants, the reactants must absorb energy from their environment to react. These reactions are endothermic and can be represented by an energy-level diagrams like Figure 1 (right).

Technically Temperature is Neither a Reactant nor Product

It is not uncommon that textbooks and instructors to consider heat as a independent "species" in a reaction. While this is rigorously incorrect because one cannot "add or remove heat" to a reaction as with species, it serves as a convenient mechanism to predict the shift of reactions with changing temperature. For example, if heat is a "reactant" (\(\Delta{H} > 0 \)), then the reaction favors the formation of products at elevated temperature. Similarly, if heat is a "product" (\(\Delta{H} < 0 \)), then the reaction favors the formation of reactants. A more accurate, and hence preferred, description is discussed below.

Exothermic and endothermic reactions can be thought of as having energy as either a "product" of the reaction or a "reactant." Exothermic reactions releases energy, so energy is a product. Endothermic reactions require energy, so energy is a reactant.

Example \(\PageIndex{1}\): Exothermic vs. Endothermic

Is each chemical reaction exothermic or endothermic?

1. \(\ce{2H2(g) + O2(g) \rightarrow 2H2O(ℓ)} + \text{135 kcal}\)
2. \(\ce{N2(g) + O2(g) + 45 kcal \rightarrow 2NO(g)}\)

Solution

No calculations are required to address this question. Just identify where the "heat" is in the chemical equation.

1. Because energy is released; this reaction is exothermic.
2. Because energy is absorbed; this reaction is endothermic.

Exercise \(\PageIndex{1}\)

If the bond energy for H-Cl is 431 kJ/mol. What is the overall bond energy of 2 moles of HCl?

Answer

Simply multiply the average bond energy of H-Cl by 2. This leaves you with 862 kJ (Table T3).

Example \(\PageIndex{2}\): Generation of Hydrogen Iodide

What is the enthalpy change for this reaction and is it endothermic or exothermic?

\[\ce{H_2(g) + I_2(g) \rightarrow 2HI(g)} \nonumber\]

Solution

First look at the equation and identify which bonds exist on in the reactants.

• one H-H bond and
• one I-I bond

Now do the same for the products

• two H-I bonds

Then identify the bond energies of these bonds from the table above:

• H-H bonds: 436 kJ/mol
• I-I bonds: 151 kJ/mol

The sum of enthalpies on the reaction side is:

436 kJ/mole + 151 kJ/mole = 587 kJ/mol.

This is how much energy is needed to break the bonds on the reactant side. Then we look at the bond formation which is on the product side:

• 2 mol H-I bonds: 297 kJ/mol

The sum of enthalpies on the product side is:

2 x 297 kJ/mol= 594 kJ/mol

This is how much energy is released when the bonds on the product side are formed. The net change of the reaction is therefore

587-594= -7 kJ/mol.

Since this is negative, the reaction is exothermic.

Example \(\PageIndex{2}\): Decomposition of Water

Using the bond energies given in the chart above, find the enthalpy change for the thermal decomposition of water:

\[\ce{ 2H_2O (g) \rightarrow 2H_2 + O_2 (g)} \nonumber \]

Is the reaction written above exothermic or endothermic? Explain.

Solution

The enthalpy change deals with breaking two mole of O-H bonds and the formation of 1 mole of O-O bonds and two moles of H-H bonds (Table T3).

• The sum of the energies required to break the bonds on the reactants side is 4 x 460 kJ/mol = 1840 kJ/mol.
• The sum of the energies released to form the bonds on the products side is
  • 2 moles of H-H bonds = 2 x 436.4 kJ/mol = 872.8 kJ/mol
  • 1 moles of O=O bond = 1 x 498.7 kJ/mil = 498.7 kJ/mol

which is an output (released) energy of 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol.

Total energy difference is 1840 kJ/mol – 1371.5 kJ/mol = 469 kJ/mol, which indicates that the reaction is endothermic and that 469 kJ of heat is needed to be supplied to carry out this reaction.

This reaction is endothermic since it requires energy to create bonds.