Energy is conserved in a chemical reaction. A chemical reaction fundamentally involves the rearrangement of atoms as reactant molecules are transformed into product molecules. Throughout this process, the total energy of the system and its surroundings remains constant. Any observable energy change, such as heat or light, is energy that has simply been converted from one form to another, not created or destroyed.
The Law of Conservation in Chemical Systems
The principle governing energy during any reaction is the Law of Conservation of Energy, also known as the First Law of Thermodynamics. This law states that energy cannot be created or destroyed; it can only be transformed from one type to another. In chemistry, the “system” is defined as the specific chemical substances undergoing the reaction.
The “surroundings” include everything else, such as the container, the air, and the laboratory environment. When a chemical change occurs, energy is exchanged between the system and the surroundings. However, the total energy contained within the combined system and surroundings must always remain constant.
If the chemical system releases energy, that same amount of energy must be absorbed by the surroundings, and vice versa. This division explains how a reaction can appear to lose or gain energy without violating the conservation law. The energy is never truly lost, only transferred across the boundary of the reacting chemicals.
Energy Storage and Exchange in Reactions
Energy is primarily stored within chemical substances as potential energy, which resides specifically in the chemical bonds holding the atoms together. This chemical potential energy results from the electrostatic forces of attraction and repulsion between the electrons and nuclei of the bonded atoms. During a chemical reaction, the process requires two distinct steps: breaking the existing bonds in the reactant molecules and forming new bonds to create the product molecules.
Breaking the initial bonds always requires an input of energy, sometimes called the bond energy. Conversely, the formation of new bonds releases energy back into the system. The net energy change observed for the entire reaction is the difference between the energy required to break the bonds and the energy released when the new, more stable bonds are formed.
If the energy released by forming the new bonds is greater than the energy required to break the old ones, the excess energy is converted and released, often as heat or light. If the opposite is true—if more energy is needed to break the bonds than is released by forming new ones—the system must absorb the difference from the surroundings. This conversion from stored chemical potential energy to other forms, such as thermal energy, is the mechanism by which energy is exchanged.
Exothermic and Endothermic Processes
The net energy exchange between the chemical system and its surroundings results in two observable types of reactions: exothermic and endothermic processes. An exothermic reaction is characterized by a net release of energy from the chemical system into the surroundings. This release manifests as heat, causing the temperature of the surroundings to increase.
A common example of an exothermic process is combustion, such as the burning of wood or natural gas. In these reactions, the chemical potential energy stored in the fuel is released as heat and light. The product molecules possess less total chemical potential energy than the original reactant molecules, and the difference is the amount transferred out to the environment.
Conversely, an endothermic reaction involves a net absorption of energy from the surroundings into the chemical system. Because the system draws energy away from its environment, the surroundings often feel cold, such as when an instant cold pack is activated. In endothermic reactions, the products have a higher chemical potential energy than the reactants, requiring a continuous supply of energy to drive the transformation.