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Unit Overview

• Unit 6: Thermochemistry

  • 6.0 Introduction - Let's Talk Energy!

  • 6.1 Endothermic and Exothermic Processes

  • 6.2 Energy Diagrams

  • 6.3 Heat Transfer and Thermal Equilibrium

  • 6.4 Heat Capacity and Calorimetry

  • 6.5 Energy of Phase Changes

  • 6.6 Introduction to Enthalpy of Reaction

  • 6.7 Bond Enthalpies

  • 6.8 Enthalpy of Formation

  • 6.9 Hess’s Law

6.0 Introduction - Let's Talk Energy!

• Forms of Energy in Thermochemistry:

  • Kinetic Energy (KE): Energy due to the motion of particles.

  • Potential Energy (PE): Energy stored in bonds of molecules.

    • Includes chemical energy and some thermal energy related to vibrations.

  • Other forms of energy include:

    • Electrical Energy: Movement of electric charges.

    • Radiant Energy: Electromagnetic energy (e.g., light, x-rays).

    • Thermal Energy: Internal energy in substances due to the motion of atoms/molecules.

    • Nuclear Energy: Energy stored in atomic nuclei.

    • Mechanical Energy & Gravitational Energy: Energy related to position or applied force.

    • Sound Energy: Movement of energy through substances in longitudinal waves.

• Key Equation for Kinetic Energy:

  [ KE = \frac{1}{2} mv^2 ]

  • Units: Joules (J), where J = kg m²/s²

• Temperature: Measure of average kinetic energy of particles.

  • Higher temperature correlates with higher average kinetic energy.

6.1 Endothermic and Exothermic Processes

• Enthalpy Change (∆H): Difference in energy between products and reactants.

  • Endothermic:

    • Absorbs energy, resulting in ∆H > 0.

    • Products have higher energy than reactants.

  • Exothermic:

    • Releases energy, resulting in ∆H < 0.

    • Reactants have higher energy than products.

6.2 Energy Diagrams

• Graphical representations showing energy changes during reactions.

  • Enthalpy change (∆H) is visualized as the difference in energy between reactants and products.

6.3 Heat Transfer and Thermal Equilibrium

• Heat Transfer: Occurs due to temperature differences until thermal equilibrium is achieved, meaning all bodies are at the same temperature.

• Average kinetic energy affects how particles move when temperature changes.

6.4 Heat Capacity and Calorimetry

• Heat Capacity (C): Quantity of energy needed to change the temperature of a substance.

  • Generally measured in J/°C.

• Specific Heat Capacity (c): Energy needed to raise the temperature of 1 g of a substance by 1 °C. For water, [ c = 4.18 \text{ J/g°C} ].

6.5 Energy of Phase Changes

• Different phases require specific amounts of energy to change states (e.g. melting, vaporizing).

• Latent Heat: Energy required for phase change without temperature change.

  • Heat of fusion (Hfus) and heat of vaporization (Hvap) are examples of latent heats.

6.6 Introduction to Enthalpy of Reaction

• Calorimetry: Technique for measuring heat changes during physical/chemical processes.

  • Uses calorimeters to determine temperature change and corresponding heat calculations.

6.7 Bond Enthalpies

• Bond Energy: Energy required to break a bond in 1 mole of gaseous molecules.

• Average bond energies give general values for bond strength.

6.8 Enthalpy of Formation

• Definition: Energy change when 1 mole of a substance is formed from elements in their standard states.

• Standard enthalpy of formation for elements is defined as zero.

6.9 Hess’s Law

• Hess's Law: States that the total enthalpy change for a reaction is the same regardless of the pathway taken.

  • Useful for calculating enthalpy changes for complex reactions.

Practice Problems

• Illustrative practice problems focus on calculating heat changes, enthalpy changes and applying the principles of thermochemistry to practical scenarios.