Chemical reactions drive all matter changes and involve a transfer of energy, fundamentally changing the energy content of the chemicals involved. Scientists classify reactions into two major categories based on the direction of this energy flow: whether the reaction releases energy into the surroundings or absorbs energy from them.
Exothermic Reactions: The Heat Producers
Exothermic reactions release energy, typically as heat or light, into the surroundings. The total energy contained in the products is lower than the energy initially held by the reactants, effectively decreasing the internal energy stored in the chemical bonds.
This energy difference is expelled, often causing the surrounding temperature to rise noticeably. When you feel warmth from a fire or a chemical hand warmer, you are directly experiencing an exothermic process. For the reaction to occur, a small amount of activation energy must first be supplied, but the energy released when new bonds form is significantly greater.
The sensation of warmth or an explosion is a direct consequence of this energy output. Exothermic processes are often self-sustaining once initiated because the released heat can continue the reaction.
Endothermic Reactions: The Heat Consumers
Endothermic reactions absorb energy from the surroundings, usually as heat. This means the total energy stored in the products is higher than the energy contained in the initial reactants. The reaction must continuously draw energy from its environment to proceed.
Because the system is pulling heat from the immediate environment, the surroundings feel cold to the touch. For instance, a chemical cold pack feels cold because the reaction inside is actively absorbing thermal energy from your skin and the air. This heat absorption causes a measurable decrease in the temperature of the reaction mixture and its environment.
In these reactions, the energy required to break reactant bonds is greater than the energy released when new product bonds form. Energy is therefore treated like a reactant, needing constant supply for the chemical change to take place. If the energy supply is cut off, the reaction will stop.
The Core Difference: Enthalpy and Temperature Change
The fundamental distinction between the two reaction types is the direction of the net energy transfer, which is formally tracked using the concept of enthalpy, denoted by the symbol Delta H. Enthalpy is a measure of the total heat content of a system at constant pressure. The sign of the change in enthalpy is the definitive marker for classification.
In an exothermic reaction, the system loses heat, which results in a negative change in enthalpy (Delta H < 0). Conversely, an endothermic reaction gains heat from the surroundings, resulting in a positive change in enthalpy (Delta H > 0). This sign convention shows whether energy is leaving or entering the chemical system.
The most observable difference is the resulting change in temperature of the surroundings. Exothermic reactions cause the temperature of the surroundings to increase because heat is released. Endothermic reactions cause the temperature of the surroundings to decrease because heat is absorbed.
Examples in Context
Exothermic reactions include all forms of combustion, such as burning wood or gasoline, which release significant heat and light. The rusting of iron, a slower process of oxidation, is also exothermic, gradually releasing a small amount of heat over time.
Hand warmers utilize the rapid oxidation of iron powder to produce heat. Another common example is the neutralization of an acid by a base, which typically releases heat and causes the mixture to warm up.
Endothermic examples include the instant cold packs used for sports injuries, where ammonium nitrate dissolves in water and rapidly absorbs heat from the environment. Photosynthesis in plants is a large-scale endothermic process, as plants absorb light energy to convert carbon dioxide and water into glucose and oxygen. The phase change of melting ice is also considered endothermic, as the ice must absorb heat from the surroundings to transition from solid to liquid water.