The study of energy and its transformations, known as thermodynamics, explains how energy moves during physical and chemical processes. Chemical reactions and physical changes always involve the exchange of energy, often in the form of heat. Enthalpy change is the specific concept scientists use to quantify this heat exchange, especially when reactions occur in open vessels exposed to the atmosphere. This measurement provides a standardized way to understand the energy balance of a process.
Defining Enthalpy and Enthalpy Change
Enthalpy, symbolized as \(H\), represents the total heat content of a system under specific conditions. Because it is impossible to measure the absolute total energy content of a substance, scientists focus on the change in enthalpy (\(\Delta H\)) rather than the absolute value of \(H\) itself.
Enthalpy change, denoted as \(\Delta H\) (read as “delta H”), is the measurable quantity that describes the heat absorbed or released during a process. The change is calculated as the difference between the final enthalpy of the products and the initial enthalpy of the reactants (\(\Delta H = H_{\text{products}} – H_{\text{reactants}}\)). This quantity has a direct meaning under the common laboratory condition of constant pressure.
When a chemical or physical process occurs at a constant pressure, the enthalpy change (\(\Delta H\)) is exactly equal to the amount of heat exchanged between the system and its surroundings. This makes \(\Delta H\) a useful tool in chemistry, as most reactions are carried out in containers open to the atmosphere where pressure remains constant. A positive \(\Delta H\) indicates the system has gained heat. Conversely, a negative \(\Delta H\) indicates the system has lost heat to the surroundings.
Identifying Energy Exchange: Exothermic and Endothermic Reactions
The sign convention of the enthalpy change determines whether a reaction releases or absorbs heat, classifying the process as either exothermic or endothermic. In an exothermic reaction, heat is released from the system into its surroundings, meaning the products have a lower enthalpy than the reactants. This release of energy is indicated by a negative value for the enthalpy change (\(\Delta H < 0[/latex]), causing the surroundings to feel warmer. A common example of an exothermic process is combustion, such as the burning of natural gas like methane, which releases a significant amount of heat energy. The chemical energy stored in the reactants is converted into heat that flows out of the system. This principle is applied in heating homes and powering engines, utilizing the released heat. Conversely, an endothermic reaction absorbs heat from its surroundings, resulting in a net gain of energy by the chemical system. Because the system gains energy, the products have a higher total enthalpy than the reactants, resulting in a positive enthalpy change ([latex]\Delta H > 0\)). When this occurs, the reaction draws heat from the immediate environment, causing the surroundings to experience a drop in temperature.
A clear, everyday example of an endothermic process is the reaction inside a chemical cold pack used for sports injuries. When the pack is activated, a salt dissolves in water and absorbs heat from the surrounding environment, making the pack feel cold to the touch. The melting of ice is also a physical change that is endothermic, as water molecules must absorb heat to break free from the solid structure.
Standard Conditions and Conceptual Measurement
Enthalpy change values are often reported under standardized conditions to ensure fair comparison between different reactions. The Standard Enthalpy of Reaction, written as \(\Delta H^\circ\), refers to the enthalpy change measured when all reactants and products are in their standard states. The standard state is defined as a pressure of 1 bar (about 100 kilopascals) and a specified temperature, commonly 298 Kelvin (25 degrees Celsius).
Scientists determine these enthalpy values through two main approaches: direct measurement and indirect calculation. Direct measurement is achieved through calorimetry, which involves measuring the temperature change of the surroundings caused by the reaction.
Direct Measurement (Calorimetry)
A device called a calorimeter, often a simple insulated container, allows researchers to calculate the heat flow. This is done by monitoring how the temperature of a known mass of surrounding substance, like water, changes.
Indirect Calculation (Hess’s Law)
For reactions too slow, too fast, or too dangerous to measure directly, scientists rely on Hess’s Law. This law states that enthalpy is a state function, meaning its change depends only on the initial and final states. Therefore, the total enthalpy change for a reaction is the same regardless of the path taken. This allows the enthalpy change of a complex reaction to be calculated by summing the known enthalpy changes of simpler, related reactions.