Combustion and Thermodynamics in Chemical Reactions

Introduction to Combustion

• Definition: Combustion is a chemical reaction involving the simplest alkane, methane (CH₄).
• Methane Details:
  • One carbon (meth) and four hydrogens.
  • Ends with "-ane" as in methane.
• Combustion Reaction:
  • Reactants: Methane and Oxygen.
  • Products: Water (H₂O) and Carbon Dioxide (CO₂).
  • Equation involves adding oxygen to produce water and carbon dioxide.

Chemical Reaction Process

Reactant Side

• Initial State:
  • Methane (CH₄) and two oxygen molecules (O₂) are low-energy reactants.
  • Reactants need energy to break into individual atoms (C, H, O).
  • Activated Complex: High-energy state where reactants are in atomic form.

Product Formation

• Transition from Activated Complex:
  • Atoms rearrange and form new bonds to create CO₂ and H₂O.
  • Energy is released when new bonds form.
• Energy in Reactions:
  • Energy added at reactant side.
  • Energy released at product side, reducing overall energy level.
  • Exothermic Reaction: Releases more energy than absorbed.

Energy Considerations in Reactions

• Exothermic Reactions:
  • Products exist at a lower energy level after energy release.
  • Feels hot as it releases energy as heat (e.g., explosions, fire).
• Endothermic Reactions:
  • More energy absorbed than released, feels cold (e.g., ice packs).

Enthalpy and Hess's Law

• Delta H (ΔH):
  • Represents the difference in energy between products and reactants.
  • Calculation: Energy of products minus energy of reactants.
  • Negative ΔH indicates exothermic; positive ΔH indicates endothermic.
• Hess's Law:
  • Principle used to calculate the enthalpy change.
  • Term "heat of reaction" sometimes used interchangeably with enthalpy.

Summary

• Combustion is a key example of exothermic reaction.
• Distinction between exothermic (heat-producing) and endothermic (heat-absorbing) reactions.
• Understanding ΔH is crucial for analyzing reaction energetics.