δu

Definition

The term δu represents the change in internal energy of a system during a process. It is a fundamental concept in thermodynamics and is crucial for understanding how energy is transferred within a system and between systems. This change in internal energy is dependent solely on the initial and final states of the system, making it a state function rather than a path function, which means that it does not depend on how the change occurs.

5 Must Know Facts For Your Next Test

1. The change in internal energy (δu) can be expressed mathematically as δu = q + w, where q is heat added to the system and w is work done on the system.
2. Unlike path functions, state functions like δu depend only on the current state of the system, meaning that different paths leading to the same endpoints will yield the same value for δu.
3. In an isolated system, δu is always zero because no heat or work can be exchanged with the surroundings.
4. For processes at constant volume, any heat added to the system contributes directly to an increase in internal energy, simplifying calculations involving δu.
5. In cyclic processes where the system returns to its original state, the total change in internal energy (δu) is also zero, reinforcing the concept that internal energy is a state function.

Review Questions

• How does δu relate to other thermodynamic quantities such as heat (q) and work (w) in a closed system?
  • In a closed system, δu is directly related to heat (q) and work (w) by the equation δu = q + w. This means that any increase in internal energy can be attributed to either heat absorbed by the system or work done on it. Understanding this relationship helps clarify how energy transfers influence the state of a system and underlines the principle that changes in internal energy occur due to various interactions within or with surroundings.
• Discuss how the concept of state functions like δu contrasts with path functions in thermodynamics.
  • State functions such as δu depend only on the current state of a system and not on how that state was reached. In contrast, path functions like work and heat depend on the specific process taken to transition between states. This distinction is crucial because it allows us to define thermodynamic properties that are consistent regardless of the process pathway, simplifying analysis in many situations. For example, no matter how much heat is exchanged or work is performed along different routes, as long as the initial and final states are unchanged, δu remains constant.
• Evaluate how understanding δu can enhance our ability to predict system behavior during thermal processes.
  • Understanding δu equips us with essential insights into how systems behave under various thermal conditions. By recognizing that δu reflects only initial and final states—independent of the process path—we can predict outcomes like energy conservation during chemical reactions or phase changes. This predictive capability is vital for applications across chemistry and engineering, enabling us to design systems that efficiently manage energy transfers while adhering to thermodynamic principles.

Related terms

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

First Law of Thermodynamics: A principle stating that energy cannot be created or destroyed, only transformed from one form to another, which relates to changes in internal energy, heat, and work.

Enthalpy: A thermodynamic quantity equal to the internal energy plus the product of pressure and volume; useful for processes occurring at constant pressure.