δu = q - w

Definition

The equation δu = q - w represents the First Law of Thermodynamics, stating that the change in internal energy (δu) of a system is equal to the heat added to the system (q) minus the work done by the system (w). This fundamental principle connects energy transfer, heat, and work in chemical processes, emphasizing that energy cannot be created or destroyed but can only change forms.

5 Must Know Facts For Your Next Test

1. The equation emphasizes the conservation of energy, showing that any change in internal energy results from heat transfer and work done.
2. In a closed system, if no heat is added (q = 0), the change in internal energy equals the negative of the work done by the system (δu = -w).
3. The sign conventions for q and w are important: heat added to the system is positive, while work done by the system is considered negative.
4. This relationship is crucial for understanding processes such as phase changes, chemical reactions, and other transformations in materials.
5. The First Law of Thermodynamics serves as a foundation for both physical chemistry and chemical engineering principles.

Review Questions

• How does δu = q - w illustrate the principle of conservation of energy within a chemical process?
  • The equation δu = q - w illustrates the principle of conservation of energy by showing that any change in a system's internal energy (δu) is directly linked to energy exchanges in the form of heat (q) and work (w). It highlights that while energy can transform from one form to another, the total amount remains constant. Therefore, if heat is added to a system or work is performed on it, its internal energy will increase, maintaining overall energy balance.
• Discuss how the sign conventions for q and w affect the interpretation of δu = q - w in practical scenarios.
  • In practical scenarios, the sign conventions for q and w play a critical role in interpreting δu = q - w. When heat is added to a system, it is considered positive, increasing internal energy. Conversely, when work is done by the system (like expanding against external pressure), it is deemed negative. This means that an increase in work output will decrease internal energy unless compensated by an input of heat. Understanding these conventions helps predict how systems will respond under different conditions.
• Evaluate how δu = q - w can be applied to analyze a specific thermodynamic process, such as combustion.
  • Applying δu = q - w to analyze combustion reveals how heat release and work impact internal energy changes. During combustion, fuel reacts with oxygen to produce heat (q) and often expand gases that perform work (w). The energy released as heat increases internal energy significantly, while expansion may lead to work being done against atmospheric pressure. By examining these exchanges, we can quantify efficiency, understand reaction conditions, and develop better energy conversion processes in chemical engineering.

Related terms

Internal Energy: The total energy contained within a system, which includes kinetic and potential energies at the molecular level.

Heat Transfer: The process of thermal energy moving from one object or system to another due to a temperature difference.

Work: The energy transferred when a force is applied over a distance, which can occur in various forms such as expansion work in gases.