Notes on Thermodynamics Lecture 1

Jul 26, 2024

Lecture Notes on Thermodynamics - Lecture 1

Introduction

  • Begin with Thermodynamics as a chapter that focuses on energy and energy transformations.
  • Clarification of difficult concepts within thermodynamics.

Basic Terms

Definition of System

  • System: The part of the universe that is under observation.
    • Examples: A person, pen, fan, tube light, or a computer can be considered systems.

Surroundings

  • Surroundings: Everything outside the system.
  • Universe: The combination of system and surroundings.

    • Formula:

      System + Surroundings = Universe

Types of Systems

Open System

  • Definition: Mass and energy can be exchanged with surroundings.
  • Example: Open pot of boiling tea.

Closed System

  • Definition: Mass cannot be exchanged; energy can be exchanged.
  • Example: A covered pot of boiling tea.

Isolated System

  • Definition: Neither mass nor energy can be exchanged.
  • Example: A thermos flask for hot or cold liquids.

Properties of Systems

Intensive Properties

  • Definition: Independent of mass and size.
  • Examples: Temperature, pressure, concentration, density, pH.

Extensive Properties

  • Definition: Dependent on mass and size.
  • Examples: Volume, mass, internal energy (U), enthalpy (H), entropy (S), free energy (G), heat capacity (C).

Important Note

  • All properties such as internal energy and enthalpy are considered state functions because they depend only on the initial and final states, irrespective of the path taken.

State and Path Functions

State Functions

  • Characteristics: Path independent, depend on the initial and final states.
  • Examples: Internal energy (U), enthalpy (H).

Path Functions

  • Characteristics: Path dependent, depend on the manner in which the transition occurs.
  • Examples: Heat (Q) and Work (W).

First Law of Thermodynamics

  • Statement: Energy cannot be created or destroyed; it can only change forms.
  • Internal energy change (╬ФU) is related to heat (Q) and work (W):

    • Formula:

      ╬ФU = Q + W

  • Sign conventions:
    • Heat absorbed by the system (Q > 0)
    • Heat released by the system (Q < 0)
    • Work done by the system (W < 0)
    • Work done on the system (W > 0)

Processes in Thermodynamics

Isobaric Process

  • Definition: Process at constant pressure.

Isochoric Process

  • Definition: Process at constant volume.

Isothermal Process

  • Definition: Process at constant temperature.

Cyclic Process

  • Definition: The initial and final states are the same; it constitutes a complete loop.

Reversible vs. Irreversible Processes

  • Reversible Process: Can be reversed without leaving any traces; proceeds through a series of equilibrium states.
  • Irreversible Process: Cannot return to the original state without changes.

Questions for Practice

  1. Identify intensive and extensive properties.
  2. Explain the significance of the first law of thermodynamics with examples.
  3. Differentiate between reversible and irreversible processes based on characteristics discussed in the lecture.

Conclusion

  • Importance of making notes for better understanding and retention of concepts discussed.
  • Recommendation to review the lecture within 24 hours for effective learning.

Full transcript