Energetics

Energetics

1. Definitions

   • Enthalpy Change
     • The overall change in energy in a reaction. Measured in KJ/mol.
   • Hess' Law
     • The energy change in a reaction is independent of the route taken.
   • Standard Enthalpy Change
     • The amount of energy required with one mole of a substance in their standard states under standard conditions.
   • Standard Conditions
     • 101kPa and 298K
   • Bond Enthalpies
     • The energy required to break a bond. This varies depending on the atoms involved and the outside conditions.

2. Equations

   • q=mcΔT
   • Enthalpy change of a reaction = energy absorbed - energy released

3. Types of Reactions

   • Endothermic Reactions
     • Takes in energy from the surroundings. Has +ve ΔH.
   • Exothermic Reactions
     • Emits energy to the surroundings. Has -ve ΔH.

4. Calorimetry

   • You need to know the mass, temperature change and the number of moles of the reactants.
   • You can just measure the temperature change of reacting solutions.
   • For combustion, place the fuel in a calorimeter and measure the temperature change.
   • Problems
     • Some heat is absorbed by the container.
     • Some heat is lost to the surroundings (this always happens - even with insulation).
     • With flammable liquids:
       • Some combustion may be incomplete.
       • Some flammable liquid may evaporate before burning.

5. General Information

   • Energy is required to break bonds and energy is released when bonds are formed.
   • Enthalpy changes include formation, combustion and reaction.
   • Formation enthalpies for all elements are zero.