Why Does Breaking a Chemical Bond Increase Potential Energy?

Breaking a chemical bond usually increases the potential energy of the atoms because energy must be supplied to separate them. The bonded arrangement is a lower-energy state than the separated atoms, much like a ball at the bottom of a valley compared with one lifted away from it.

This describes an individual bond. A whole chemical reaction can still release energy if forming new bonds releases more energy than breaking the original bonds requires.

Potential energy along a bond

For two atoms, potential energy depends on the distance between their nuclei. When the atoms are far apart, they interact only weakly. As they approach, attractive interactions lower the potential energy until it reaches a minimum at the bond’s equilibrium length. If they are pushed closer than that, strong repulsion makes the potential energy rise sharply.

To break the bond, the atoms must be pulled from the minimum out to a large separation. The potential energy rises along the way, approaching the energy of the separated atoms. The energy difference between the minimum and this separated-atom limit is the bond’s dissociation energy for the specified molecular state and definition.

In a potential-energy diagram, the bound molecule is therefore in a well. The depth of the well indicates how much energy is needed to separate the atoms, subject to distinctions such as whether vibrational zero-point energy is included.

Why separating atoms takes energy

At the equilibrium bond length, attractive and repulsive effects balance. Moving the atoms apart from that position requires work against the net attractive force. That work becomes additional energy of the system. In thermochemical language, a specified dissociation is endothermic when its enthalpy change is positive; at constant pressure under the usual conditions, this corresponds to heat absorption.

The energy does not have to be supplied as heat. It may come from light, electrical energy, collisions, or energy released elsewhere in a reaction. These are ways of supplying energy; “endothermic” specifically refers to a positive enthalpy change and, at constant pressure under the usual conditions, heat absorption.

For a simple gas-phase dissociation, the bond dissociation enthalpy is the energy required for a specified homolytic cleavage under stated conditions. Homolytic cleavage gives one electron from the bond to each atom. Bond dissociation enthalpy is closely related to the increase in potential energy, but it is not always exactly the same number: enthalpy includes thermal contributions, while a potential-energy curve describes molecular energy as a function of geometry. For a precise comparison, use the same physical definition and conditions throughout.

Worked example: separating H₂

Consider the gas-phase process

$$\mathrm{H_2(g) \rightarrow 2H(g)}$$

The H–H bond dissociation enthalpy is about $436\ \mathrm{kJ\,mol^{-1}}$. So, under the usual gas-phase bond-enthalpy convention, roughly $436\ \mathrm{kJ}$ must be supplied for each mole of H₂ molecules dissociated. The two hydrogen atoms are at a higher energy than the bonded H₂ molecule, with the energy difference accounted for by the input.

On an electronic potential-energy curve, the well depth is often denoted $D_e$ and is measured from the bottom of the well to the separated-atom limit. The molecule also has vibrational zero-point energy, so the energy needed to dissociate from its lowest vibrational level is smaller than $D_e$. This is why different, carefully defined dissociation-energy values need not match exactly.

Breaking bonds versus the energy of a reaction

A reaction’s energy change depends on both bonds broken and bonds formed. Breaking bonds costs energy; forming bonds generally releases energy because the products’ bonded atoms move to lower potential energy.

A useful estimate for a gas-phase reaction is

$$\Delta H_{\mathrm{reaction}} \approx \sum D(\text{bonds broken}) - \sum D(\text{bonds formed})$$

Here, $D$ values are bond dissociation energies. The estimate is approximate when it uses average bond energies, because a bond’s strength depends on its specific molecular environment.

For example, if breaking reactant bonds requires $700\ \mathrm{kJ\,mol^{-1}}$ and forming product bonds releases $900\ \mathrm{kJ\,mol^{-1}}$, the estimate is

$$\Delta H_{\mathrm{reaction}} \approx 700 - 900 = -200\ \mathrm{kJ\,mol^{-1}}$$

The reaction is exothermic overall even though its bond-breaking steps require energy. Energy released during bond formation more than compensates for the energy absorbed during bond breaking.

Common mistakes

• Saying that breaking a bond releases energy. Separating the bonded atoms requires energy. The mistaken claim often comes from confusing bond breaking with the overall energy change of a reaction.
• Assuming every reaction that breaks bonds is endothermic. A reaction can break some bonds and still release energy overall if it forms stronger or more numerous product bonds.
• Treating potential energy, internal energy, and enthalpy as identical. They are related but distinct quantities. A bond’s potential-energy well, a spectroscopic dissociation energy, and a measured bond dissociation enthalpy can differ because their definitions and conditions differ.
• Thinking “stored energy” means energy is released when the bond is broken. A bonded state is generally lower in potential energy than separated atoms. Energy must be added to reach the separated state.