Bond Energies

Atoms bond together to form compounds because they attain lower energies in this state compared to when they are individual atoms. A quantity of energy, equal to the difference between the bonded atoms' energies and those of the separated atoms, is released, typically as heat. In essence, the bonded atoms possess lower energy than individual atoms do. When atoms combine to create a compound, energy is released, and the compound has an overall lower energy.

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When a chemical reaction occurs, molecular bonds break, and other bonds form to create different molecules. For instance, the bonds of two water molecules break to form hydrogen and oxygen.

\[ 2H_2O \rightarrow 2H_2 + O_2\]

Energy is always required to break a bond, known as bond energy. Although the concept might seem straightforward, bond energy serves a crucial role in explaining the structure and properties of a molecule. It can help determine the most appropriate Lewis Dot Structure when multiple Lewis Dot Structures are possible.

Energy is always required to break a bond. Energy is released when a bond is made.

While each molecule has a characteristic bond energy, we can make some generalizations. For instance, although the exact value of a C-H bond energy depends on the particular molecule, all C-H bonds have a bond energy of roughly the same value because they are all C-H bonds. It takes about 100 kcal of energy to break 1 mol of C-H bonds, so we say that the bond energy of a C-H bond is around 100 kcal/mol. A C-C bond has an approximate bond energy of 80 kcal/mol, and a C=C bond has about 145 kcal/mol. We can calculate a more general bond energy by averaging the bond energies of a specific bond in different molecules.

When a bond is strong, there is a higher bond energy because more energy is needed to break a strong bond. This correlates with bond order and bond length. When the bond order is higher, bond length is shorter. The shorter the bond length, the greater the bond energy due to increased electric attraction. Generally, the shorter the bond length, the greater the bond energy.

The average bond energies in Table T3 are the averages of bond dissociation energies. For example, the average bond energy of an O-H bond in H₂O is 464 kJ/mol. This arises because the H-OH bond requires 498.7 kJ/mol to dissociate, while the O-H bond needs 428 kJ/mol.

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\[\dfrac{498.7\; kJ/mol + 428\; kJ/mol}{2}=464\; kJ/mol\]

Considering more bond energies of the bond in different molecules will result in a more accurate average. However,

• Average bond values are not as accurate as molecule-specific bond-dissociation energies.
• Double bonds are higher energy bonds compared to single bonds (but not necessarily 2-fold higher).
• Triple bonds have even higher energy than double and single bonds (but not necessarily 3-fold higher).

Bond Breakage and Formation

When a chemical reaction occurs, the atoms in the reactants rearrange their chemical bonds to form products. This new arrangement of bonds doesn't have the same total energy as the bonds in the reactants. Hence, a chemical reaction always results in an energy change.

In some reactions, the energy of the products is lower than the reactants. Thus, during the reaction, substances lose energy to the surroundings. These reactions are exothermic and can be depicted by an energy-level diagram as in Figure 1 (left). In most cases, the energy is released as heat (though some reactions emit energy as light). In reactions where the products have higher energy than the reactants, the reactants must absorb energy from their environment to react. These reactions are endothermic and can be illustrated by diagrams like Figure 1 (right).

Technically Temperature is Neither a Reactant nor Product

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

Exothermic and endothermic reactions can be viewed as having energy as either a "product" or a "reactant." Exothermic reactions release energy, making energy a product. Endothermic reactions require energy, making energy a reactant.

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

Is each chemical reaction exothermic or endothermic?

1. \(2H_{2(g)} + O_{2(g)} \rightarrow 2H_2O_{(&#;)} + \text{135 kcal}\)
2. \(N_{2(g)} + O_{2(g)} + \text{45 kcal} \rightarrow 2NO_{(g)}\)

Solution

No calculations are required to address this question. Just look at where the "heat" is in the chemical reaction.

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 gives you 862 kJ (Table T3).

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