How Bonds Are Broken and Formed in Chemical Reactions
In chemical reactions, atoms rearrange as bonds in the reactants break and new bonds form in the products. Breaking bonds requires energy; forming bonds releases energy. The reaction’s approximate enthalpy change is:
ΔH ≈ total energy required to break bonds − total energy released when bonds form
A negative ΔH indicates an exothermic reaction, which releases energy to its surroundings. A positive ΔH indicates an endothermic reaction, which absorbs energy. This energy balance does not, by itself, determine how quickly the reaction occurs.
Breaking bonds requires energy; forming bonds releases it
Separating bonded atoms takes an input of energy. When atoms form a bond, energy is released as the bonded arrangement forms.
The overall energy change depends on both processes. A reaction can release energy even though breaking its reactant bonds requires energy: this happens when forming the product bonds releases more energy than breaking the reactant bonds takes. If breaking the reactant bonds takes more energy than forming the product bonds releases, the reaction absorbs energy.
Example: methane combustion
The balanced equation for methane combustion is:
CH₄ + 2 O₂ → CO₂ + 2 H₂O
The reaction breaks the C–H bonds in methane and the O=O bonds in oxygen. It forms C=O bonds in carbon dioxide and O–H bonds in water. Because forming the product bonds releases more energy than breaking the reactant bonds requires, methane combustion is exothermic.
A bond-energy calculation must count every bond in the balanced equation, including the coefficients. For example, the two water molecules contain four O–H bonds in total:
ΔH ≈ total bond energy of bonds broken − total bond energy of bonds formed
Bond energies are average values for particular bond types, so this method estimates rather than gives an exact value for a specific reaction. Measured thermochemical data are more suitable when greater precision is needed. The NIST Chemistry WebBook provides data for many substances.
For more on the method and enthalpy of reaction, see Chemistry LibreTexts: Bond Energies and Chemistry LibreTexts: Enthalpy of Reaction. Khan Academy: Bond Energy and Enthalpy also walks through bond-energy calculations.
Energy change does not determine reaction speed
A reaction’s energy change and its rate are separate properties. An exothermic reaction can proceed slowly if it faces a large activation-energy barrier—the initial energy needed to follow a reaction pathway. A spark can provide enough energy to start methane combustion; after ignition, the reaction releases energy as product bonds form.
A catalyst can provide a pathway with a lower activation-energy barrier. It changes how readily a reaction proceeds, but not its reactants or products, so it does not change the reaction’s overall enthalpy change.
Endothermic reactions also occur. They absorb energy from their surroundings or from an external source, such as heat or light. The bond-energy balance describes the reaction’s energy change, but the pathway and conditions also affect whether and how quickly it proceeds.
Key Takeaways
- Breaking bonds requires energy; forming bonds releases energy.
- Approximate ΔH equals the energy required to break reactant bonds minus the energy released when product bonds form.
- Negative ΔH indicates an exothermic reaction; positive ΔH indicates an endothermic reaction.
- Bond-energy calculations use average values and provide estimates; measured thermochemical data are better for precise values.
- Reaction energy and reaction speed are different. Activation energy affects the rate, while a catalyst can lower the activation-energy barrier without changing overall enthalpy.
