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
Universe = System + Surroundings
First Law of Thermodynamics: ΔE_universe = 0
ΔE = q + w
Enthalpy Change: ΔH = q_p
Hess's Law: Heat of reaction is fixed and independent of the path.
Adjust enthalpy values to account for balanced equations.
Bond energies can provide insights into enthalpy changes.