General Terminology

• Bonds/attractive forces:

  • Breaking bonds/attractive forces requires energy (energetically unfavorable, positive change).
  • Forming bonds/attractive forces releases energy (energetically favorable, negative change).

• System: part of the universe where bonds are broken/formed at a molecular or ionic level (part of universe that is studied).

• Potential energy diagrams/reaction energy profile:

  

  Potential energy diagrams.

  • Potential energy (kJ/mol) vs. reaction coordinate (reaction progress).
  • Activation energy barrier: minimal energy for a reaction to occur; the energy difference between the activated complex (peak; split second where bonds transition from breaking to forming) and the reactants.
    • The height of the activation energy determines the rate of the reaction.
  • A: endothermic reaction (enthalpy is positive).
  • B: exothermic reaction (enthalpy is negative).
  • Lower potential energy means the compound is more stable.

Heat

<aside> 🌡️ Heat: transfer of energy between two bodies at different temperatures that proceeds until equilibrium. Units are J or kJ.

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• Change in heat energy is represented by q.
• The value of q changes depending on the amount of reactants.

Endothermic & Exothermic Reactions

• The direction sign (+/-) is always from the system perspective.
• Chemical or physical processes can be endothermic/exothermic.
  • Example of chemical changes: combustion of ethanol.
  • Example of physical change: vaporization of ethanol.

<aside> 🌡️ Endothermic: positive enthalpy ($+\Delta H^o$), net input of energy into the system.

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• Requires more energy to break bonds than energy released by bond formation.
• Heat absorbed/input/consumed.

<aside> 🌡️ Exothermic: negative enthalpy ($-\Delta H^o$), net output of energy from the system.

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• Releases more energy than required to break bonds.
• Heat evolved/produced/formed.
• More thermodynamically favorable (lowering potential energy of the system).

Enthalpy ($\Delta H$)