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The is the cornerstone of in physical systems. It states that energy can't be created or destroyed, only converted between forms. This principle is crucial for understanding how energy flows in various processes.

Applying the First Law helps us analyze different thermodynamic processes like isothermal, adiabatic, and isobaric changes. By understanding these processes, we can predict how systems behave and calculate important quantities like work, heat, and internal energy changes.

First Law of Thermodynamics

Application of first law of thermodynamics

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  • Fundamental principle states energy cannot be created or destroyed, only converted from one form to another (law of )
  • For a , change in internal energy (ΔU\Delta U) equals heat added to the system (QQ) minus work done by the system (WW)
    • Mathematical representation: ΔU=Q−W\Delta U = Q - W
  • Heat (QQ) is positive when added to the system (heating) and negative when removed from the system (cooling)
  • Work (WW) is positive when done by the system (expansion) and negative when done on the system (compression)
  • Calculate changes in internal energy by:
    1. Determining heat added to or removed from the system (calorimetry)
    2. Calculating work done by or on the system (pressure-volume diagrams)
    3. Using the equation to find ΔU\Delta U

Relationships in thermodynamic processes

  • Isothermal process: temperature remains constant
    • Change in internal energy is zero (ΔU=0\Delta U = 0)
    • Heat added equals work done by the system (Q=WQ = W)
      • Example: isothermal expansion of an ideal gas
  • Adiabatic process: no heat exchanged with surroundings
    • Change in internal energy equals negative of work done (ΔU=−W\Delta U = -W)
      • Example: rapid compression of a gas in an insulated cylinder
  • Isobaric process: pressure remains constant
    • Heat added equals change in internal energy plus work done by the system (Q=ΔU+WQ = \Delta U + W)
      • Example: heating a gas at constant pressure
  • Isochoric (isovolumetric) process: volume remains constant
    • No work is done (W=0W = 0)
    • Heat added equals change in internal energy (Q=ΔUQ = \Delta U)
      • Example: heating a gas in a rigid container
  • : system remains infinitesimally close to equilibrium throughout the entire process

State functions vs path-dependent quantities

  • depend only on the current state, not the path taken to reach that state
    • Examples: internal energy (UU), enthalpy (HH), entropy (SS)
    • Changes in state functions (ΔU\Delta U, ΔH\Delta H, ΔS\Delta S) are independent of the path between initial and final states
      • Example: change in internal energy is the same for isothermal and adiabatic processes between the same initial and final states
  • depend on the specific path taken between initial and final states
    • Examples: heat (QQ) and work (WW)
    • Values of heat and work vary depending on the process or path, even if initial and final states are the same
      • Example: work done in an isothermal expansion differs from work done in an adiabatic expansion between the same initial and final volumes
  • Identify state functions and path-dependent quantities to apply appropriate equations and principles when analyzing thermodynamic processes

Types of Thermodynamic Processes

  • : can be reversed without leaving any change in the system or surroundings
  • : cannot be reversed without leaving a change in the system or surroundings
  • : system returns to its initial state after a series of changes
  • : system is in thermal, mechanical, and chemical equilibrium with its surroundings

Key Terms to Review (16)

Closed System: A closed system is a thermodynamic system that does not exchange matter, but can exchange energy, with its surroundings. It is an idealized model used in the study of thermodynamics, particularly in the context of the First Law of Thermodynamics.
Conservation of Energy: The conservation of energy is a fundamental principle in physics which states that the total energy of an isolated system remains constant, it is said to be conserved over time. Energy can neither be created nor destroyed; rather, it can only be transformed or transferred from one form to another.
Cyclic Process: A cyclic process is a thermodynamic process in which a system undergoes a series of changes and eventually returns to its initial state, completing a cycle. This type of process is fundamental to the understanding of the First Law of Thermodynamics, as it allows for the analysis of the relationships between work, heat, and the internal energy of a system.
Energy conservation: Energy conservation is the principle stating that the total energy of an isolated system remains constant. It implies that energy can neither be created nor destroyed, only transformed from one form to another.
First law of thermodynamics: The First Law of Thermodynamics states that energy cannot be created or destroyed in an isolated system, only transformed from one form to another. It is also known as the law of energy conservation.
First Law of Thermodynamics: The First Law of Thermodynamics states that energy can be transformed from one form to another, but it cannot be created or destroyed. It establishes the fundamental principle of energy conservation, which is crucial for understanding heat transfer, thermodynamic systems, and adiabatic processes in an ideal gas.
Irreversible Process: An irreversible process is a thermodynamic process that cannot be reversed, meaning the system and its surroundings cannot be returned to their initial states without causing some permanent change. This concept is fundamental to the understanding of the First Law of Thermodynamics and the concept of entropy.
Isovolumetric process: An isovolumetric process, also known as an isochoric process, is a thermodynamic process in which the volume of a system remains constant while the pressure and temperature may change. This type of process is significant because it illustrates how energy can be transferred within a system without any change in volume, allowing for the exploration of energy conservation and transformation.
Path-dependent quantities: Path-dependent quantities are properties that depend on the specific path taken during a process rather than just the initial and final states. In thermodynamics, this concept is crucial because it highlights how energy transfer and work can vary based on the route followed during a thermodynamic cycle or transformation.
Quasistatic Process: A quasistatic process is a thermodynamic process that occurs slowly and in near-equilibrium conditions, allowing the system to remain infinitesimally close to equilibrium at all times. This type of process is often used in the analysis of the First Law of Thermodynamics, as it allows for the clear distinction between the work done on or by the system and the heat transferred to or from the system.
Reversible process: A reversible process is a thermodynamic process that can be reversed without leaving any net change in either the system or the surroundings. These processes are ideal and occur infinitesimally slowly, allowing the system to remain in equilibrium throughout.
Reversible Process: A reversible process is a thermodynamic process that can be reversed without leaving any trace on the surroundings. In other words, a reversible process can be undone, and the system and the surroundings can be returned to their initial states without the expenditure of any work or the absorption of any heat from the surroundings.
State function: A state function is a property of a system that depends only on its current state, not on the path taken to reach that state. This means that when considering changes in the system, only the initial and final states matter, not the process or steps taken to get from one to the other. Understanding state functions is crucial because they help simplify the analysis of thermodynamic processes, allowing us to focus on measurable properties rather than complex histories.
State functions: State functions are properties of a system that depend only on the current state of the system, not on the path taken to reach that state. Examples include internal energy, pressure, and temperature.
Thermodynamic Equilibrium: Thermodynamic equilibrium is a state in which the macroscopic properties of a system, such as temperature, pressure, and chemical composition, do not change over time. It is a fundamental concept in thermodynamics that underpins the study of energy transformations, work, heat, and the behavior of systems.
ΔU = Q - W: The equation ΔU = Q - W represents the First Law of Thermodynamics, which states that the change in internal energy (ΔU) of a system is equal to the work (W) done on the system minus the heat (Q) transferred to the system. This fundamental relationship describes the conservation of energy and the interconversion between different forms of energy.
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Glossary
Glossary
3.4 Thermodynamic Processes