Thermodynamics: Overview

Thermodynamic Equilibrium

Definition: A state where there is no spontaneous change in a system's conditions.
Example: Gas in a cylinder at uniform temperature and pressure, balanced by an external force.
Processes:
- Change occurs only with external influence on state functions (temperature, volume).
- Reversible Process: System is near equilibrium at each step, changes can be reversed.
- Irreversible Process: Sudden changes, like a balloon bursting, lead to non-equilibrium states.

Concept: Fundamental to thermodynamics but complex to define precisely.
Measurement: Temperature is assessed when two objects reach thermal equilibrium.
- Involves the Zeroth Law of Thermodynamics.
Temperature Scales:
- Celsius Scale: 0°C (freezing point), 100°C (boiling point) at 1 atm.
- Fahrenheit Scale: 32°F (freezing point), 212°F (boiling point).
- Absolute Scales:
  - Kelvin (K): Related to Celsius, K = C + 273.15.
  - Rankine (R): Related to Fahrenheit, R = F + 459.67, R = 1.8 K.
- Absolute Zero: Zero point of Kelvin and Rankine scales.

Energy: In physics, it has precise but complex meanings, often related to work.
Work: Product of force and displacement.
- No work done without movement.
Kinetic and Potential Energy:
- Kinetic: Energy of motion.
- Potential: Energy of position; conversion between forms in conservative systems.
Conservation of Energy: Total energy constant in closed systems; relates to the First Law of Thermodynamics.
Heat: Different from other energy forms; conversion of work to heat is not fully reversible.
Second Law of Thermodynamics: Limits conversion of heat back to mechanical energy.

Microscopic Understanding:
- Increase in kinetic energy of molecules due to added thermal energy.
- Includes rotational, vibrational, and chemical energy within molecules.
Constitutes:
- Total internal energy = sum of all energy forms in a thermodynamic state.
- Total system energy includes internal energy plus kinetic and potential energy due to motion and elevation.