LESSON_2_Q4_THERMOCHEMISTRY

Thermochemistry Overview

• Thermochemistry deals with the heat changes involved in chemical reactions and physical processes.

Learning Points

• Key concepts to cover:

  • Conservation of energy in chemical reactions.

  • Definitions of exothermic and endothermic reactions.

  • Introduction to thermodynamics.

Energy Changes in Chemical Reactions

• Changes in matter typically involve energy changes:

  • Endothermic process: Energy is absorbed.

  • Exothermic process: Energy is released.

  • Bond breaking requires energy input while bonding releases energy.

System vs. Surroundings

• System: The specific part of the universe being studied.

• Surroundings: Everything else in contact with the system.

• Energy (usually in the form of heat) can be exchanged between system and surroundings.

  • In endothermic processes, heat flows into the system.

  • In exothermic processes, heat flows out to the surroundings.

Types of Systems

• Open system: Exchanges both matter and energy.

• Closed system: Exhanges only energy.

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

Classification of Reactions

• Identify processes as exothermic or endothermic:

  1. Melting of butter: Endothermic

  2. Rubbing hands with alcohol: Endothermic

  3. Burning gasoline: Exothermic

  4. Mixing HCl with water: Exothermic

  5. Subliming naphthalene: Endothermic

  6. Perspiring: Endothermic

  7. Making of ice: Exothermic

  8. Ice forming in clouds: Exothermic

  9. Inflating a bicycle tire: Exothermic

  10. Breaking down food: Exothermic

First Law of Thermodynamics

• Thermodynamics: Study of energy exchanges.

• Energy defined as the ability to perform work or transfer heat.

• Internal energy (E): Total of all kinetic and potential energy in a system.

• The first law states:

  • ∆E = q + w

    • q: heat added to the system (positive for endothermic)

    • w: work done on the system (positive for work done on the system)

  • For an exothermic process, both q and w are negative.

Internal Energy Changes

• Change in internal energy

  • Example: If the system absorbs 523 J of heat and does 452 J of work:

    • Given: q = +523 J, w = -452 J

    • ∆E = q + w = 523 J - 452 J = 71 J

  • Result: Internal energy increases by 71 J.

Enthalpy Change

• Enthalpy (H): Heat transfer at constant pressure.

• Change in enthalpy (ΔH) indicated in thermochemical equations.

  • If ΔH > 0 (+), process is endothermic.

  • If ΔH < 0 (-), process is exothermic.

• Enthalpy is a state function; depends on the quantity and physical state of reactants/products.

Rules for Thermochemical Equations

1. ΔH is proportional to mass - if coefficients double, ΔH doubles.

2. Reaction has equal and opposite ΔH for reverse reactions.

Examples of Enthalpy Changes

1. Combustion of methane: Exothermic; ΔH is negative.

2. Burning of sulfur: Also exothermic; ΔH is negative.

3. Melting of ice: Endothermic; requires energy (ΔH is positive).

Enthalpy Changes in Physical Processes

Key Types:

• Enthalpy of solution: Heat absorbed/released when solute dissolves (e.g., NaOH in water).

• Enthalpy of fusion: Energy required to melt a solid (e.g., ice to water).

• Enthalpy of vaporization: Energy required to convert a liquid to gas (e.g., water to steam).

• Enthalpy of neutralization: Heat change during acid-base neutralization.

• Enthalpy of condensation: Heat absorbed when a gas becomes a liquid.

Activities and Calculations

• Practice calculating enthalpy changes:

  • Example: Calculate ΔH when 38.0 g of water evaporates.

  • Given: ΔHvap = 40.9 kJ/mol.

Conclusion

• Understanding thermochemistry is crucial for analyzing heat changes in chemical processes and their real-world applications.

• Equipped with foundational thermodynamics concepts, students can tackle complex reaction mechanisms and calculations.