Daley_Student_Abbrev_Unit_I_2025_Chem152

Thermodynamics I: Energy

Chapter Overview

• Focus on the system, surroundings, and the universe.

Kinetic and Potential Energy

• Kinetic Energy: Energy of motion or energy that is being transferred.

  • Thermal Energy: A type of kinetic energy associated with temperature.

• Potential Energy: Energy stored in an object or energy associated with an object's composition and position.

  • Chemical Energy: A type of potential energy due to atomic structure, attachments, and positions relative to each other.

• Energy can only be transferred, not created or destroyed.

First Law of Thermodynamics

• The energy of the universe is constant: ∆E_universe = 0 Joules.

• First Law of Thermodynamics: Law of Conservation of Energy states that total energy is unchanging; it can only change forms (absorbed or released) but cannot be created or destroyed.

• The energy change of a system (∆E_system) plus the change of its surroundings (∆E_surroundings) must equal 0 J.

• Equation: ∆E_universe = ∆E_system + ∆E_surroundings

  • Energy release from the system is equal to energy absorption by surroundings, and vice versa.

Understanding Thermodynamic Systems

• System: Matter and associated energy under study (reaction mixture).

• Surroundings: Everything else (flask, room, universe).

• The energy can flow between the system and surroundings:

  • System drops in energy = Surroundings gain energy.

  • System gains energy = Surroundings lose energy.

Energy Terms and Signs

• Magnitude: Proportional to the size of the energy value; larger values indicate more energy.

• Absorption or Emission: Indicated by the signs:

  • Positive (+) indicates energy is absorbed into the system.

  • Negative (–) indicates energy is emitted from the system.

Heat and Work

• Heat (q): Thermal energy exchanged with surroundings, measured in joules (J).

• Work (w): Energy used to move an object against a force; depends on force magnitude and distance.

• Work done on or by the system:

  • w_surroundings = – w_system

State Functions vs. Path Functions

• State Function: Describes a specific state or change in states, independent of the path taken.

• Path Function: Dependent on the process taken to move between states.

• Internal Energy: The sum of the kinetic and potential energies of all particles in a system.

  • Change in internal energy (∆E) is given by ∆E = E_final – E_initial.

Energy Diagrams and Reaction Enthalpy

• Energy is a state function; heat and work are path functions.

• Example Reaction: CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)

• The change in internal energy relates to the energy of the products and reactants.

Quick Summary of Internal Energy

• Internal energy analogy to a bank account:

  • Withdrawals (energy out) indicated by negative signs.

  • Deposits (energy in) indicated by positive signs.

• If reactants have higher internal energy than products, ΔE_system is negative (energy flows out).

Gas Expansion and Contraction Work

• Gas Expansion Work: System does work on surroundings, energy/volume changes cause negative w_system.

• Gas Contraction Work: Surrounds do work on the system, resulting in positive w_system.

Specific Heat and Heat Transfer

• Heat transfer from hot to cold materials until thermal equilibrium is reached:

  m_metal × C_s,metal × ΔT_metal = − (m_water × C_s,water × ΔT_water)

• Relationship defined through calorimetry.

Quantitating Heat and Work

• Heat (q) related through:

  • q = C × ΔT (where C is heat capacity).

  • Specific heat (C_s) indicates energy needed to raise temp of 1 g substance by 1 °C.

Measuring ΔE

• ΔE can be determined from q + w where work can equal zero at constant volume.

• In practice, differences in surrounding temperatures are observed, typically utilizing bomb calorimetry for measurements.

Thermochemical Concepts

• Enthalpy (H): Sum of internal energy and PV (pressure-volume) product.

• ΔH of a reaction signifies heat at constant pressure.

• ΔH can generally replace q in ΔE = q + w.

Heat of Formation Reactions

• Formation reactions create 1 mole of products from elemental forms at 25°C.

• Standard heat of formation denoted as ΔH_f°.

Enthalpy Calculation Examples

• Utilize ΔH values from balanced chemical equations to determine reaction heat changes.

• Calculate based on stoichiometry of reactions and standardized values for formation heats.

Review of Key Thermochemistry Points

1. Universe = System + Surroundings

2. First Law of Thermodynamics: ΔE_universe = 0

3. ΔE = q + w

4. Enthalpy Change: ΔH = q_p

5. Hess's Law: Heat of reaction is fixed and independent of the path.

6. Adjust enthalpy values to account for balanced equations.

7. Bond energies can provide insights into enthalpy changes.