Energy Relationships in Chemical Reactions

Energy Concepts

  • Energy: The capacity to do work.
    • Radiant Energy: Comes from the sun; primary energy source of the Earth.
    • Thermal Energy: Energy associated with the random motion of atoms/molecules.
    • Chemical Energy: Energy stored in the bonds of chemical substances.
    • Nuclear Energy: Energy within protons and neutrons in an atom.
    • Potential Energy: Energy available by virtue of an object's position.

Heat Transfer

  • Heat: Transfer of thermal energy between two bodies at different temperatures.
  • Temperature: Measure of thermal energy.
  • Formula: Temperature = Thermal Energy.

Thermochemistry

  • Thermochemistry: Study of heat change in chemical reactions.
  • System: Specific part of the universe under study.
    • Open System: Exchange of mass & energy.
    • Closed System: Exchange of energy only.
    • Isolated System: No exchange (neither mass nor energy).

Energy Changes in Reactions

  • Exothermic Process: Releases heat; thermal energy transferred from the system to surroundings.
    • Example: [ 2H2 (g) + O2 (g) \rightarrow 2H_2O (l) + ext{energy} ]
  • Endothermic Process: Absorbs heat; heat supplied to the system from surroundings.
    • Example: [ H2O (g) + ext{energy} \rightarrow H2O (l) ]

Thermodynamic Principles

  • Thermodynamics: Study of heat and energy interconversion.
  • State Functions: Properties depend on the state of the system regardless of how the state was achieved (e.g., energy, pressure, volume, temperature).

First Law of Thermodynamics

  • First Law: Energy can be converted but not created or destroyed.
    • Formula: ( ext{D}U{system} + ext{D}U{surroundings} = 0 ) or ( ext{D}U{system} = - ext{D}U{surroundings} )
  • Chemical reaction example: [ S(s) + O2 (g) \rightarrow SO2 (g) ] (Exothermic)

Internal Energy Changes

  • Internal Energy Change: ( ext{DU} = q + w )
    • Where ( q ): heat exchange, ( w ): work done.
    • Work done on/by the system: ( w = -P ext{D}V )

Work Done by Gases

  • Work expression: ( w = F \times d )
  • For gas expansion: ( w = -P ext{D}V )
  • Distinction: Positive work indicates compression; negative work implies expansion.

Heat Calculations in Reactions

  1. Heat Capacity: ( C = m \times s ) (mass × specific heat)
  2. Heat Change Formula: ( q = m \times s \times \Delta T )
    • ( ext{D}T = T{final} - T{initial} )

Enthalpy

  • Enthalpy (H): Quantifies heat flow into/out of a system at constant pressure.
    • Enthalpy change formula: ( ext{D}H = ext{H}{products} - ext{H}{reactants} )
    • For exothermic ( ext{D}H < 0 ) (heat released) and for endothermic ( ext{D}H > 0 ) (heat absorbed).

Thermochemical Equations

  • Example: [ H2O (s) \rightarrow H2O (l), \text{D}H = 6.01 \, kJ/mol ] (Absorbs heat: Endothermic)
  • Reversing reaction changes sign of ( ext{D}H ).
  • Coefficients represent moles; ( 2H2O (s) \rightarrow 2H2O (l), \text{D}H = 2 \times 6.01 = 12.0 \, kJ ).

Heat of Reaction

  • Directly related to the amount of substance reacted and its enthalpy change. Use stoichiometry for calculations.

Hess's Law

  • Enthalpy change for a reaction is the same regardless of how it's carried out—important in calculating unknown enthalpies.

Standard Enthalpy of Formation

  • Defined as heat change when 1 mol of a compound forms from its elements under standard conditions.
    • Always set to zero for elements in their most stable form (e.g., ( ext{D}H^0 (O_2) = 0 )).

Summary of Key Formulas

  • ( ext{D}U = q + w )
  • ( ext{D}H = q ) at constant pressure
  • ( q = m \times s imes \Delta T )