When a chemical reaction occurs, energy is always transferred, usually as heat, leading to a noticeable temperature change. This transfer is governed by the law of conservation of energy, meaning energy is merely moved between substances, not created or destroyed. Understanding the direction of heat flow dictates whether a reaction heats up or cools down. This energy movement determines the final temperature of the reaction and its surroundings, providing a clear classification for the process.
Defining the Reaction System and Surroundings
To properly track energy transfer, scientists divide the universe into two parts: the system and the surroundings. The system is the specific part of the universe being studied, referring to the reacting chemicals themselves. The surroundings include everything else outside the system, such as the container, the air, and a thermometer used to measure the change.
Heat flow is always analyzed from the perspective of the chemical system. If the system gains energy, the reaction is classified one way, and if it loses energy, it is classified the opposite way. Crucially, any temperature change we measure is a change in the surroundings, since the thermometer is part of the surroundings. A temperature increase in the surroundings means the system has transferred heat out, and a temperature decrease means the system has absorbed heat.
Exothermic Processes: The Source of Temperature Increase
A temperature increase in the surroundings is the definitive sign of an exothermic process. An exothermic reaction releases energy, typically as heat, from the chemical system into the surrounding environment. This transfer of thermal energy makes the container feel hot to the touch and causes the temperature reading to rise.
The source of this released energy lies in the chemical bonds of the reacting molecules, which contain stored potential energy. For a reaction to proceed, reactant bonds must first be broken, requiring an energy input. New bonds are then formed to create the products, and this formation always releases energy. In an exothermic reaction, the energy released during product bond formation is greater than the energy required to break the reactant bonds.
This net excess energy is expelled from the system as heat, leading to the measured temperature increase in the surroundings. A common example is the combustion of a fuel like natural gas, where the formation of new, more stable bonds releases a large amount of heat. This principle is also used in chemical hand warmers, where the oxidation of iron powder releases heat. Dissolving compounds like calcium chloride in water also generates heat, which is why it is used for de-icing roads.
Endothermic Processes: Absorbing Heat
An endothermic process involves the absorption of energy, usually heat, by the chemical system from its surroundings. This absorption of thermal energy causes the measured temperature of the surroundings to drop, making the reaction container feel cold. The system is drawing heat away from its environment to fuel the chemical change.
In endothermic reactions, the energy required to break the bonds in the reactants is greater than the energy released when the new product bonds are formed. This energy deficit must be made up by drawing heat from the surroundings, which results in the cooling effect. The products of an endothermic reaction therefore contain more stored potential energy than the original reactants did.
A practical illustration of this process is an instant cold pack, which often contains ammonium nitrate and water separated by a membrane. When the membrane is broken, the ammonium nitrate dissolves in the water, absorbing a significant amount of heat from the surrounding environment. Another large-scale natural example is photosynthesis, where plants absorb light energy to convert carbon dioxide and water into glucose and oxygen. Even the simple act of ice melting is endothermic, as the ice absorbs heat to break the bonds holding the water molecules in a solid structure.