1733815493_Dina ch 15 section 3 Thermochemical equations ---- (1)

Chapter 15: Energy and Chemical Change

Section 15.3: Thermochemical Equations

  • Main Idea: Thermochemical equations express the amount of heat released or absorbed during chemical reactions.

Let’s Engage

  • Exercises:

    • Compare endothermic and exothermic reactions regarding changes in enthalpy.

    • Explain the six phase changes in terms of heat absorption and release.

    • Define a thermochemical equation and compare it to a chemical equation.

    • Identify conditions where heat (q) equals change in enthalpy (ΔH).

Comparison of Reactions

Endothermic Reactions

  1. Definition: Absorb energy from surroundings (heat).

  2. Change in Enthalpy (∆H): Positive (+∆H), energy absorbed to break bonds.

  3. Example: Melting ice, photosynthesis.

Exothermic Reactions

  1. Definition: Release energy to surroundings (heat).

  2. Change in Enthalpy (∆H): Negative (−∆H), energy released when new bonds are formed.

  3. Example: Combustion of fuels, freezing water.

Phase Changes and Heat Transfer

Six Phase Changes:

  1. Melting (Solid → Liquid):

    • Heat absorbed.

    • Example: Ice melting into water.

  2. Freezing (Liquid → Solid):

    • Heat released.

    • Example: Water freezing into ice.

  3. Vaporization (Liquid → Gas):

    • Heat absorbed.

    • Example: Water boiling into steam.

  4. Condensation (Gas → Liquid):

    • Heat released.

    • Example: Water vapor condensing into liquid water.

  5. Sublimation (Solid → Gas):

    • Heat absorbed.

    • Example: Dry ice sublimating into carbon dioxide.

  6. Deposition (Gas → Solid):

    • Heat released.

    • Example: Frost forming from water vapor.

Thermochemical Equations

Definition

  • A thermochemical equation includes balanced chemical equations with physical states of reactants/products and energy changes represented by the change in enthalpy.

Differences from Chemical Equations

  • Energy Information:

    • Chemical Equation: Shows only reactants and products.

    • Thermochemical Equation: Includes energy changes (ΔH).

    • Examples:

      • Chemical Equation: H2 + O2 → H2O

      • Thermochemical Equation: H2 + 1/2O2 → H2O, ΔH = −286 kJ/mol

Conditions of Heat and Enthalpy

  • Heat (q) equals ΔH when reactions occur under constant pressure.

    • Description: At constant pressure, heat exchanged is related directly to the enthalpy change.

Writing Thermochemical Equations

  • Two Methods:

    1. ΔH at the End:

      • Example: 2NO2(g) → N2(g) + 2O2(g), ΔH = −67.4 kJ

    2. Energy as Reactant/Product:

      • If endothermic, write energy as reactant;

      • If exothermic, write energy as product.

    • Ex: SO2(g) + 297 kJ → S(s) + O2(g) 2NO2(g) → N2(g) + 2O2(g) + 67.4 kJ

Enthalpy of Combustion

  • Defined as the enthalpy change for the complete burning of one mole of a substance.

  • Includes standard enthalpies of combustion that vary by substance.

Understanding Phase Changes

Energy Changes and States

  • Endothermic Processes: heat absorbed; results in vaporization and melting.

  • Exothermic Processes: heat released; results in condensation and freezing.

  • The heat involved in vaporization and fusion is numerically equal but has opposite signs for condensation and solidification.

Practical Applications

  • Farmers may use the heat of fusion of water to protect crops by causing water to freeze and release heat into the surrounding air, preventing frost damage.

Independent Practice Problems

Exercises on Phase Changes and Combustion

  • Define and analyze phase changes based on heat absorption or release.

  • Explore combustion reactions to understand the conceptual links between enthalpy changes and observable reactions.

Biological Connection

  • The energy released from glucose metabolism parallels its combustion, highlighting biological processes relying on combustion reactions to obtain cellular energy (ATP).

Recap of Combustion Reactions

  1. Combustion Definition: Reaction of a fuel with oxygen, releasing energy.

  2. Examples:

    • Fuel (hydrocarbon) + O2 → CO2 + H2O (exothermic)

    • Common fuels: methane (CH4), octane (C8H18).

Review and Summary

  • Thermochemical equations are crucial for understanding energy exchanges in chemical reactions and physical processes, leading to better comprehension of thermodynamics in both practical and biological contexts.