CHEM_18_Module_5 (2024-2025) handout

Module 5: Chemical Thermodynamics

  • Thermodynamics: Study of the transformations of energy

    • Origin of the term: from Greek "therme" (heat) and "dynamis" (power).

Types of Thermodynamics

  • Chemical Thermodynamics: Energy changes during chemical reactions.

  • Chemical Kinetics: Rates of reactions and mechanisms.

Types of Energy

  • Potential Energy (PE):

    • Definition: Stored energy due to position or composition.

    • Sources: Forces holding atoms in a nucleus, chemical bonds, and Intermolecular Forces of Attraction (IMFA).

  • Kinetic Energy (KE):

    • Definition: Energy of motion.

    • Types of motions at molecular level:

      • Translational: Movement from one location to another.

      • Vibrational: Oscillations of atoms in a molecule.

      • Rotational: Spinning of molecules.

Units of Energy

  • Calorie (cal): Amount of energy to raise 1 g of water by 1°C.

  • Joule (J): SI unit of energy.

    • Conversion: 1 cal = 4.184 J.

Systems vs. Surroundings

  • System: Part of the universe of interest in a study.

  • Surroundings: The rest of the universe.

Types of Systems

  • Open Systems: Exchange of matter and energy with surroundings

  • Closed Systems: Exchange of energy only.

  • Isolated Systems: No exchange of energy or matter.

Internal Energy (U) of a System

  • Total energy within a system: U = KE + PE.

  • Absolute value of U is unmeasurable; only changes in Internal energy (ΔU) are significant.

  • Heat (Q): Energy transfer between system and surroundings, stopping when thermal equilibrium is reached.

  • Work (W): Energy transfer involving physical force, often represented by: W = -PΔV for pressure-volume work.

First Law of Thermodynamics

  • Law of Conservation of Energy:

    • Energy can be converted but cannot be created/destroyed.

    • Formula: ΔU = Q + W.

Energy Transformation

  • When an object moves from a higher potential energy (PE) position to a lower PE position, the energy converts to kinetic energy (KE).

Spontaneous Reactions

  • Reactions that occur without external intervention.

  • Driven to lower potential energy states; examples include rusting, melting ice, and water flow.

Gibbs Free Energy (G)

  • Unifies parameters like enthalpy (ΔH) and entropy (ΔS) for predicting reaction spontaneity.

    • Formula: ΔG = ΔH - TΔS.

Calorimetry

  • The process of measuring heat released or absorbed during reactions.

  • Based on the conservation of energy: heat lost by the system is gained by surroundings.

  • For phase changes at constant pressure:

    • Q = nΔH (where n = moles).

Phase Changes

  • Endothermic: Energy absorbed, leading to melting and boiling.

  • Exothermic: Energy released, represents freezing and condensation.

Phase Diagrams

  • Graphs illustrating the states of a substance at varying temperatures and pressures.

  • Critical points indicate conditions beyond which a gas cannot become a liquid.

Practice Problems

  • Identify system types for scenarios (e.g., coffee in a cup vs. a helium-filled balloon).

  • Calculate heat evolved in reactions under different conditions.