Thermochemistry Notes

Thermochemistry Overview

Section 6.1: Energy and Energy Units

  • Energy Definition:

    • The capacity to do work or transfer heat.

  • Types of Energy:

    • Potential Energy (PE): Energy possessed by an object due to its position, composition, or condition.

    • Kinetic Energy (KE): Energy that an object possesses due to its motion.

Energy Units

  • SI Unit of Energy: Joule (J)

    • Conversion: $1 ext{ J} = 5 imes 10^2 ext{ kg m}^2/ ext{s}^2$

  • Calorie (cal):

    • Defined as the amount of energy needed to raise the temperature of 1 g of water by 1℃.

    • Conversion: $1 ext{ cal} = 4.184 ext{ J}$

Section 6.2: Energy, Heat, and Work

  • Definitions:

    • Work (w): Energy resulting from a force acting on an object over a distance.

    • Heat (q): Transfer of energy due to a temperature difference.

  • Systems:

    • Open System: Matter and energy can exchange with surroundings.

    • Closed System: Only energy can exchange.

    • Isolated System: No exchange of matter or energy occurs.

  • Conservation of Energy:

    • Energy cannot be created or destroyed, only transformed.

    • Internal Energy Change: riangle E = q + w

Section 6.3: Energy as a State Function

  • State Functions:

    • Properties dependent only on the state of the system (e.g., internal energy).

  • Path Functions:

    • Depend on the path taken to reach a particular state (e.g., work and heat).

Section 6.4: Energy and Enthalpy

  • Enthalpy (H): A state function that combines internal energy (U), pressure (P), and volume (V).

    • Change in Enthalpy: riangle H = q_P (heat flow at constant pressure)

  • Pressure-Volume Work:

    • Given by the formula: w = -P riangle V

Endothermic and Exothermic Reactions

  • Endothermic Reactions: Heat flows into the system.

  • Exothermic Reactions: Energy flows out of the system.

Section 6.5: Specific Heat

  • Definition of Specific Heat (c): Energy required to raise the temperature of 1 g of a substance by 1℃.

    • Common units: J/(g·℃).

  • Heat Flow Equation:

    • q = m imes c imes riangle t where m = mass, c = specific heat capacity, riangle t = change in temperature.

Section 6.6: Calorimetry

  • Calorimetry:

    • The study of heat transfers by measuring temperature changes in involved substances.

  • Constant-Pressure Calorimetry: Typically done in a Styrofoam cup; measures heat change at atmospheric pressure, denoted q_P = riangle H.

  • Constant-Volume Calorimetry (Bomb Calorimeter): Measures energy changes when substances combust at constant volume.

Section 6.7: Enthalpy in Chemical Reactions

  • Enthalpy Change (ΔH):

    • Equal to the heat transfer at constant pressure.

    • Independent of the path taken (state function).

  • Hess’s Law:

    • The total enthalpy change is the sum of the changes for the individual steps of a reaction.

    • Rules for Manipulating Chemical Equations:

    1. Reversing an equation changes the sign of its ΔH.

    2. Multiplying coefficients alters the enthalpy value by the same factor.

    3. Summing reactions results in the sum of their enthalpy changes.

Section 6.8: Standard Enthalpies of Formation

  • Definition: The enthalpy change when 1 mol of a compound forms from its elements in their standard state (25°C, 1 atm).

    • Standard enthalpy of formation for elements in their standard state is 0.

  • Examples of Standard Enthalpies of Formation:

    • Methane (CH4) = -74.6 kJ/mol

    • Carbon Dioxide (CO2) = -393.5 kJ/mol

    • Water (H2O) = -285.8 kJ/mol

Calculation Examples:

  • Specific Heat Calculation: Using q = m imes c imes riangle t.

  • Combustion of Decane in Bomb Calorimeter: Measure heat changes using total heat capacity: q = C imes riangle t.

Overall, these notes cover fundamental concepts in thermochemistry, focusing on the definitions of energy forms, the relations between heat, work, and systems, as well as the laws governing energy and enthalpy in chemical reactions.