Strike a match and you feel heat on your fingers. Snap an instant cold pack during a sports game and it suddenly becomes icy against your skin. Both involve chemical reactions, yet one warms you and the other cools you. What controls whether a reaction releases energy to the surroundings or absorbs it? The key lies in how much energy is stored in the chemical bonds of the reactants and products, and how those total bond energies change during the reaction.
Chemical reactions are all around you, and many of them clearly involve energy changes:
In each of these examples, energy is either flowing out of the reacting system into the surroundings (heating the surroundings) or flowing into the system from the surroundings (cooling the surroundings). To understand why, we need to look more closely at how atoms and bonds store and exchange energy.
Matter is made of atoms, and atoms themselves have structure. Each atom has a tiny positively charged nucleus containing protons and neutrons, surrounded by negatively charged electrons. At the atomic scale, forces between electric charges control how atoms interact:
When atoms approach each other, the attraction between the positively charged nuclei of one atom and the negatively charged electrons of another can create a stable arrangement we call a chemical bond. Forming a bond lowers the potential energy of the system because the oppositely charged particles are in a more favorable arrangement.
Chemical reactions re-arrange which atoms are bonded to which. That means they change the pattern of electric-charge interactions at the atomic scale, and therefore change how much energy is stored in the system. This change in stored energy shows up as energy being released to or absorbed from the surroundings.
To connect energy changes in reactions to bond changes, chemists use the idea of bond energy. Bond energy is the amount of energy required to break a particular chemical bond between two atoms in one mole of molecules, in the gas phase. Strong bonds (like the bond in \(\textrm{N}_2\)) require a lot of energy to break; weak bonds require less.
In any chemical reaction, some bonds in the reactant molecules are being broken, and new bonds in the product molecules are being formed. The total energy change for the reaction depends on the combination of all these changes. When thinking qualitatively, it is helpful to imagine adding up all the energies needed to break the original bonds and comparing that to all the energies released when new bonds form, as suggested in [Figure 1].
Two key ideas:
The total bond energy of the reactants (considering all their bonds) and the total bond energy of the products are generally not the same. The difference between these totals determines whether the reaction as a whole releases or absorbs energy, as suggested in [Figure 1].
As you compare different reactions, you can think: if forming the new product bonds releases more energy than it took to break the original bonds, then overall energy will be released. If breaking the original bonds costs more energy than is released from forming new bonds, then overall energy will be absorbed.
We can now build a conceptual model that connects bond changes to energy flow without doing any numerical calculations. Consider a general reaction:
\[ \textrm{Reactants} \rightarrow \textrm{Products} \]
Our model has three main steps:
The overall energy change of the reaction depends on the balance between the energy needed in step 2 and the energy released in step 3. Our model can be summarized qualitatively as:
\[ \textrm{Energy change} \approx \textrm{energy to break bonds} - \textrm{energy released forming bonds} \]
If the energy released when product bonds form is greater than the energy required to break reactant bonds, the reaction releases energy to the surroundings. If the energy required to break reactant bonds is greater than the energy released by forming product bonds, the reaction absorbs energy from the surroundings.
This model emphasizes that the release or absorption of energy depends on the changes in total bond energy between reactants and products.
When a reaction releases energy to the surroundings, it is called exothermic. In exothermic reactions, the products end up with lower total chemical potential energy than the reactants, even though the individual bonds in the products are often stronger (and therefore have higher bond energies). The “excess” energy is released, often as heat, sometimes also as light or other forms.
In terms of our bond-based model, exothermic reactions have this pattern:
Common examples include:
In each case, the fact that the surroundings get warmer is evidence that the reaction system has moved to a lower-energy state, with products that have stronger or more stable bonding overall than the reactants.
On an energy diagram, exothermic reactions show products at a lower energy level than reactants. This connects back to total bond energy: lower final energy corresponds to product bonds that collectively represent a lower potential energy configuration.
When a reaction absorbs energy from the surroundings, it is called endothermic. In endothermic reactions, the products have higher total chemical potential energy than the reactants. This extra stored energy comes from the surroundings, often as heat, making the surroundings cooler.
In terms of bonds:
Examples include:
In these reactions, the final bonded arrangements store more energy per mole than the initial ones. Our model explains this: the total energy required to break all necessary bonds in the reactants is larger than the total energy released from forming new bonds in the products.
On an energy diagram, endothermic reactions show products at a higher energy level than reactants, representing the higher total bond energy of the products.
Energy diagrams are a powerful way to visualize how total bond energy changes during a reaction. They plot the potential energy of the system against the progress of the reaction from reactants to products, as illustrated in [Figure 2]. These diagrams do not show individual bonds, but they summarize the overall energy story of the reaction.
Two main features appear on most reaction energy diagrams, as summarized in [Figure 2]:
For an exothermic reaction:
For an endothermic reaction:
The activation energy “hump” exists in both cases because energy must be put in initially to break or weaken existing bonds before new ones can form.
Thinking back to bond energies, the vertical distances in these diagrams reflect how the total bond energy of the products compares to that of the reactants. When we say the products are at lower energy, we are saying that the set of bonds in the products corresponds to a more stable, lower potential energy arrangement of charges than in the reactants.
Later, when analyzing real-world technologies like batteries or fuel cells, scientists and engineers use these diagrams and related ideas to reason about how much energy a reaction will release or absorb without necessarily performing detailed bond-energy calculations at every step.
At the deepest level, bond energies and reaction energy changes come from interactions between electric charges. When two atoms approach each other, the positively charged nuclei and negatively charged electrons interact in complex ways.
There are three main types of interactions to consider between two atoms:
At a certain distance, the attractive forces balance the repulsive forces, and the potential energy of the system is at a minimum. This distance corresponds to the bond length, and the depth of this energy “well” represents the bond strength: a deeper well means a stronger bond with higher bond energy (more energy needed to pull the atoms apart).
Different kinds of bonds arise from different ways charges interact, as summarized in [Figure 3]:
Stronger attractions (for example, between ions with larger charges or atoms with strongly overlapping electron clouds) create deeper potential energy wells and thus higher bond energies. When a reaction replaces a set of relatively weak bonds with a set of stronger bonds, the system falls into a deeper energy well overall and releases energy to the surroundings. The opposite is true when stronger bonds are replaced by weaker ones; the system climbs to a higher energy level and must absorb energy.
So, at the atomic scale, the release or absorption of energy in a chemical reaction can be traced directly back to rearrangements of electric charges and the resulting changes in potential energy.
Using our bond-energy model, we can compare reactions qualitatively without doing numerical calculations.
Example comparison 1: Combustion vs cold pack dissolution
Example comparison 2: Photosynthesis vs cellular respiration
In both comparisons, the key is to focus on the change in total bond energy from reactants to products. We do not need to calculate the exact amounts to reason about whether energy is released or absorbed.
These examples also connect back to the shapes of energy diagrams in [Figure 2] and the idea of stronger bonds corresponding to deeper energy wells as in [Figure 3].
Understanding how total bond energies change in reactions is crucial in many real-world applications and engineering designs 💡:
In all these cases, scientists and engineers rely on models of bond energies and reaction energy diagrams to reason about how much energy will be released or absorbed and under what conditions, without always calculating detailed bond-energy sums for every reaction.
• Chemical reactions involve breaking old bonds in reactants and forming new bonds in products. Breaking bonds always absorbs energy; forming bonds always releases energy.
• Bond energy is a measure of how much energy is required to break a bond. Stronger bonds correspond to deeper potential energy wells at the atomic scale, caused by stronger electrostatic attractions between nuclei and electrons.
• The overall energy change of a reaction depends on the change in total bond energy from reactants to products. If forming product bonds releases more energy than is needed to break reactant bonds, the reaction is exothermic and releases energy. If more energy is needed to break reactant bonds than is released by forming product bonds, the reaction is endothermic and absorbs energy.
• Energy diagrams model these changes by showing reactants, products, and the activation energy hump. Exothermic reactions have products at lower energy than reactants; endothermic reactions have products at higher energy.
• At the atomic scale, changes in bond energies arise from rearrangements of electric charges and the balance of attractive and repulsive forces between nuclei and electrons.
• Real-world technologies such as fuels, batteries, thermal packs, and biological processes rely on carefully chosen reactions where the change in total bond energy produces the desired energy release or absorption.
Together, these ideas form a coherent model that explains why the release or absorption of energy from a chemical reaction system depends on how the total bond energy changes when reactants are transformed into products.
