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Chapter 6 - CHEM 1113 Broering

Chapter 6: Thermochemistry

Review: Energy and Energy Units

  • Energy: Capacity to do work or transfer heat.

  • Heat (q): Form of energy that flows due to temperature differences.

  • Kinetic Energy (KE): Defined by the equation KE = ½ mv²

  • Units of Energy:

    • Joule (J): SI unit of energy (1 J = 1 kg·m²/s²)

    • Calorie (cal): Energy needed to raise 1 g of water by 1°C (1 cal = 4.184 J)

    • Kilocalories: 1 kcal = 1000 cal (nutritional Calorie)

Systems in Thermochemistry

  • System Definition: Can be a galaxy, laboratory apparatus, or simply particles involved.

  • Types of Systems:

    • Isolated System: No matter or energy transfer.

    • Closed System: Energy can flow but not matter.

    • Open System: Both energy and matter can flow freely.

Laws of Thermodynamics

  • Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

  • First Law of Thermodynamics: Total energy of the universe is constant (E_universe = 0).

    • Changes in the system are mirrored in the surroundings: ΔE_system = -ΔE_surroundings.

Internal Energy

  • Internal Energy (U): Sum of all kinetic and potential energies of particles within a system.

  • Change in Internal Energy (ΔU):

    • Calculated as ΔU = U_final - U_initial.

    • ΔU > 0 indicates an increase, ΔU < 0 indicates a decrease.

Heat and Work in Energy Transfer

  • Ways Energy is Transferred:

    • Heat (q): Always from hot to cold objects.

    • Work (w): Energy transferred due to a force moving an object (w = F × d).

  • Internal Energy Change: ΔU = q + w.

Heat Transfer: Exothermic and Endothermic Processes

  • Exothermic (−q): Releases energy from the system to the surroundings.

  • Endothermic (+q): Absorbs energy from surroundings into the system.

Pressure-Volume Work (PV Work)

  • Work done during gas expansion/compression.

  • For constant pressure: w = -PΔV.

State Functions and Path Functions

  • State Functions: Independent of the path taken (e.g., Internal Energy, Temperature).

  • Path Functions: Depend on the path (e.g., Work, Heat).

Enthalpy

  • Enthalpy (H): A state function related to internal energy, pressure, and volume.

  • At constant pressure, ΔH = ΔU + PΔV.

  • Change in enthalpy (ΔH) equals heat transferred at constant pressure (q_p).

Specific Heat and Heat Capacity

  • Heat Capacity (C or C_P): Energy needed to raise temperature of an object by 1°C at constant pressure.

  • Specific Heat (c): Energy required to raise temperature of 1 g of a substance by 1°C.

  • Molar Heat Capacity (c_n): Energy required to raise the temperature of 1 mole of a substance by 1°C.

Calorimetry: Measuring Energy Changes

  • Calorimetry: Technique to measure heat changes in a chemical process.

  • Constant-Pressure Calorimetry: Miamimal heat transfer between system and surroundings, used for reactions in solution at constant pressure.

Example of Calorimetry

  • Heat lost by metal equals heat gained by water.

    • Used for specific heat calculations of different materials.

Hess’s Law

  • Hess’s Law: Enthalpy (ΔH) for a process equals the sum of ΔH for individual steps.

  • Applicable to manipulate thermochemical equations.

Standard Enthalpies of Reaction

  • Use enthalpy of formation data to determine ΔH for reactions.

  • Standard conditions are 1 atm pressure and 25°C; ΔH°f (standard enthalpy of formation) is defined under these conditions.

Chapter 6: Thermochemistry

  • Energy: Capacity to do work or transfer heat. Heat (q) flows due to temperature differences.

  • Kinetic Energy (KE): KE = ½ mv².

  • Energy Units: Joule (J) = kg·m²/s², Calorie (cal) = energy to raise 1 g water by 1°C (1 cal = 4.184 J), Kilocalories (1 kcal = 1000 cal).

  • System Definitions:

    • Isolated: No transfer of matter or energy.

    • Closed: Energy transfer possible, no matter transfer.

    • Open: Both energy and matter can transfer.

  • Laws of Thermodynamics: Energy cannot be created or destroyed; total energy in the universe is constant. Changes in system reflect in surroundings (ΔE_system = -ΔE_surroundings).

  • Internal Energy (U): Total kinetic and potential energies; change (ΔU) calculated as ΔU = U_final - U_initial.

  • Energy Transfer: Heat (q) flows hot to cold; work (w) done by force (w = F × d). ΔU = q + w.

  • Heat Transfer: Exothermic (−q) releases energy; Endothermic (+q) absorbs energy.

  • PV Work: Work done in gas expansion/compression, w = -PΔV for constant pressure.

  • State vs. Path Functions: State functions are path-independent (e.g., Internal Energy); path functions depend on the process (e.g., Work, Heat).

  • Enthalpy (H): ΔH = ΔU + PΔV at constant pressure, equals heat transfer (q_p).

  • Specific Heat: Heat capacity (C_P) needed to raise an object's temperature; specific heat (c) = energy to raise 1 g by 1°C; molar heat capacity (c_n) applies to 1 mole.

  • Calorimetry: Technique for measuring heat changes, particularly in constant-pressure settings. Heat lost by one substance equals heat gained by another.

  • Hess’s Law: Total ΔH = sum of ΔH for individual steps; used for thermochemical equations.

  • Standard Enthalpies: ΔH can be calculated using enthalpy of formation data at standard conditions (1 atm, 25°C).