Comprehensive Thermodynamics and System Properties Notes

Definitions and Primary Concepts of Thermodynamics

• Thermodynamics is defined as the branch of science dealing with the study of heat changes in all types of physical and chemical processes.
• The entire universe is divided into two primary parts:     $\text{Universe} = \text{System} + \text{Surrounding}$
• System:
  • The specified portion of the universe chosen for thermodynamic study.
• Surroundings:
  • The remaining portion of the universe outside the system.
• Boundary:
  • Systems and surroundings are separated by either a physical boundary or an imaginary boundary wall.
  • Diathermic Boundary Wall: A boundary wall that is a good conductor of heat, allowing thermal transfer between the system and surroundings.
  • Adiabatic Boundary Wall: A boundary wall that is a bad or poor conductor of heat, preventing thermal exchange between the system and surroundings.

Classification of Thermodynamic Systems

• Systems are categorized based on their ability to exchange energy and matter with their surroundings:
  • Open System:
    • A system that can exchange both matter and energy with its surroundings.
    • Examples: Hot coffee in an open cup, living beings, and $NaCl$ in an open test tube.
  • Closed System:
    • A system that can exchange energy with its surroundings, but cannot exchange matter.
    • Examples: Hot tea in a sealed stainless steel vessel or cup.
  • Isolated System:
    • A system that can exchange neither energy nor matter with its surroundings.
    • Examples: Hot tea inside a thermoflask.
    • The Universe as an Isolated System: The entire universe is considered an isolated system where the total energy remains constant (e.g., energy distributions such as $100\,J$, $20\,J$, $50\,J$, and $10\,J$ sum up to a constant total), which illustrates the First Law of Thermodynamics.

Phase-Based Classification of Systems

• Homogeneous Systems:
  • A system that is uniform throughout and consists of a single phase.
  • Examples: Pure solids, pure liquids, gaseous mixtures, and true solutions.
• Heterogeneous Systems:
  • A system that is not uniform throughout and consists of more than one phase.
  • Examples:
    • Solid in contact with liquid.
    • Liquid in contact with liquid (immiscible liquids).
    • Solid in contact with gas or vapour.
    • Colloidal solutions (where Solute = dispersed phase and Solvent = dispersion medium).
    • Suspensions.

Macroscopic Systems and Macroscopic Properties

• Macroscopic Systems:
  • A system made up of a very large number of particles such as atoms, molecules, or ions (e.g., quantities on the order of $10^4$ to $10^{10}$ particles, or macroscopic quantities like $0.00189 \times N_A \times 3$ atoms).
• Macroscopic Properties:
  • Properties associated with a macroscopic system, classified into two major categories: Intensive Properties and Extensive Properties.

Intensive Properties

• Macroscopic properties that do not depend on the amount, quantity, or bulk of the substrate present in the system.
• Comprehensive Examples:
  • Temperature
  • Pressure
  • Density
  • Melting Point (MP)
  • Boiling Point (BP)
  • Freezing Point (FP)
  • Surface Tension
  • Molar Volume
  • Viscosity
  • Refractive Index
  • Molar Concentration
  • Specific Heat
  • Latent Heat
  • $\text{pH}$
  • Electromotive Force (EMF) / Cell Potential

Extensive Properties

• Macroscopic properties that depend on the amount, quantity, or bulk of the substrate present in the system.
• Comprehensive Examples:
  • Mass
  • Volume
  • Internal Energy ($U$ or $E$)
  • Enthalpy ($H$)
  • Entropy ($S$)
  • Free Energy ($G$)
  • Heat ($Q$)
  • Work ($W$)

Mathematical Relationships Between Property Types

• Ratio of Two Extensive Properties:     $\frac{\text{Extensive Property}}{\text{Extensive Property}} = \text{Intensive Property}$
  • Example:         $\frac{\text{Mass (Extensive)}}{\text{Volume (Extensive)}} = \text{Density (Intensive)}$
• Product of Intensive and Extensive Properties:     $\text{Intensive Property} \times \text{Extensive Property} = \text{Extensive Property}$

State Functions and Path Functions

State Functions

• Macroscopic properties that depend exclusively on the final and initial states of the system, independent of the intermediate path or steps taken to reach that state.
• Absolute values of state functions cannot be directly determined or calculated; only changes in state functions are measured:     $\Delta \text{State Function} = \text{Final State} - \text{Initial State}$
• For a chemical reaction:     $\Delta \text{Property} = \text{Property}_{\text{product}} - \text{Property}_{\text{reactant}}$
• Key Examples:
  • Enthalpy ($H$): Change in enthalpy is $\Delta H = H_p - H_r = H_{\text{product}} - H_{\text{reactant}}$
  • Free Energy ($G$): Change in free energy is $\Delta G$
  • Entropy ($S$): Change in entropy is $\Delta S$
  • Internal Energy ($U$): Change in internal energy is $\Delta U$

Path Functions

• Properties that depend on the specific path or intermediate steps taken during a thermodynamic process.
• Key Examples:
  • Work ($W$)
  • Heat ($Q$)
• Special Conversion Condition:
  • Heat ($Q$) transferred at constant pressure ($Q_p$) is equal to Enthalpy ($H$), which converts heat into a state function equivalent under constant pressure conditions.

Types of Thermodynamic Processes

• Isothermal Process:
  • A thermodynamic process in which temperature remains constant.
• Adiabatic Process:
  • A thermodynamic process in which no heat exchange occurs between system and surroundings.
• Isochoric Process:
  • A thermodynamic process occurring at constant volume.
• Isobaric Process:
  • A thermodynamic process occurring at constant pressure.