Chapter 6 New 20

Thermochemistry and Thermodynamics

  • Thermochemistry: Study of heat changes in chemical reactions.

  • Thermodynamics: Study of interconversion of heat and other energies.

  • Energy: Ability to do work (Work = Force × Distance).

  • Types of Energy:

    • Kinetic Energy (KE) = ½ mv² (energy of motion).

    • Potential Energy (stored energy based on position).

  • Conservation of Energy: Kinetic and potential energy are interconvertible; energy cannot be created or destroyed.

Types of Energies

  • Radiant Energy: From the sun (primary energy source).

  • Thermal Energy: Associated with random motion of atoms and molecules.

  • Chemical Energy: Energy stored in chemical bonds.

  • Nuclear Energy: Energy from neutrons and protons in an atom.

Study of the System

  • Chemistry focuses on system (reaction of interest) versus surroundings.

  • Heat flow is essential: Heat flows, cold does not.

  • Classify reactions as exothermic (heat released) or endothermic (heat absorbed).

Types of Systems

  • Open System: Exchange of mass and energy.

  • Closed System: Exchange of heat only.

  • Isolated System: No exchange of heat or mass.

State Functions

  • Governed by thermodynamic laws; include macroscopic properties (volume, energy, pressure, temperature).

  • Path independent; depend only on initial and final conditions.

ΔU = U final – U initial (Internal energy change).

First Law of Thermodynamics

  • Energy conservation: Energy cannot be created or destroyed, only transferred.

  • Mathematically: ΔU system = -ΔU surroundings.

Energy, Work, and Heat

  • Energy: Capability to do work and transfer heat: ΔU = q + w (q = heat, w = work).

  • Definitions:

    • -q: heats absorbed (endothermic).

    • -w: work done on the system (positive).

  • q and w are path dependent, not state functions.

Determining Work and Heat

  • Work: w = F × d (mechanical work, often in gas expansion/compression).

  • Expansion: w = -PΔV (P = external pressure).

  • Gas expands: ΔV positive, w negative.

  • Gas contracts: ΔV negative, w positive.

Enthalpy

  • At constant volume: ΔU = qv (no PV work).

  • At constant pressure: ΔU = qp - PΔV; Enthalpy: H = U + PV; ΔH = ΔU + PΔV.

  • Negative ΔH: Exothermic reaction; Positive ΔH: Endothermic reaction.

  • ΔHrxn = H(products) – H(reactants).

Stoichiometry of Thermochemical Reactions

  • Heat changes associated with specific mole amounts.

  • Reaction wording matters for enthalpy sign.

  • For reversed reactions: ΔH sign reverses.

Standard Enthalpy of Formation

  • Heat change for formation of one mole of a compound from elements at standard conditions.

  • Elements in stable form: ΔH = 0.

Hess’s Law

  • Total enthalpy change is path independent; use for calculating standard enthalpy of reactions.

Specific Heat and Heat Capacity

  • Specific Heat (s): Heat required to raise 1g of substance by 1°C.

  • Heat Capacity (C): Heat required to raise a specific quantity (mass).

  • q = msΔT; q = CΔT.

Heats of Solution and Dilution

  • Heat of solution: Heat change when a solute dissolves.

  • Avoid adding water to concentrated solutions due to exothermic reactions.

Problem Solving and Reflection

  • Understand energy changes and thermochemical principles.

  • Recognize importance in real-life examples such as exothermic reactions (e.g., combustion) and endothermic reactions (e.g., photosynthesis).