Why do we usually use ΔH^∘ rather than ΔE^∘ when we discuss energies of reaction? Develop an arg

step 1 When we discuss energies during reactions, most of the time we talk about those in terms of enth

Why do we usually use ΔH^∘ rather than ΔE^∘ when we discuss...

Key Concepts

Enthalpy (?H)

Enthalpy is a thermodynamic quantity equivalent to the total heat content of a system and is defined as the sum of the internal energy and the product of its pressure and volume. It is especially useful in chemical reactions carried out at constant pressure because the heat absorbed or released in such processes is directly related to ?H, making it easier to measure and interpret experimentally.

Internal Energy (?E)

Internal energy represents the total energy contained within a system, including all kinetic and potential energies at the molecular level. While ?E provides a complete picture of the energy changes within a system, it does not account for the work done by or against the system due to expansion or compression, which is why it is less commonly used in constant pressure conditions.

Pressure-Volume Work

Pressure-volume work is the energy involved when a system expands or contracts, quantified by the product of pressure and volume change (p?V). In many chemical reactions, particularly those involving gases, this work is significant because the enthalpy change, ?H, includes both the internal energy change and the work done against atmospheric pressure, linking ?H closely to measurable heat flow.

Experimental Conditions (Constant Pressure vs. Constant Volume)

The selection of ?H or ?E hinges on the experimental conditions of the reaction. Most reactions are conducted at constant pressure, which means that the heat exchanged is equal to the change in enthalpy. However, for systems at constant volume—such as in a bomb calorimeter—the change in internal energy, ?E, is the more directly observed quantity, suggesting a rationale for its use in certain contexts.

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