TOPICS COVERED:

ENTHALPY

• Energy = heat
  • Energy transferred = heat transferred
  • Actually heat and work but mostly heat (work negligible, chemical reactions are inefficient)
• Definition of energy: the ability to do work
• Types of energy:
  • Kinetic energy
  • Potential energy
• Units of energy:
  • Joules (J)
  • Kilojoules (kJ)
  • Calories (cal)
    • 1 cal = 4.184 J
      • Amount of heat needed to raise the 1 gram (1 mL) of water 1ºC
  • Kilocalories (kcal) or food Calories (Cal)
• Heat = enthalpy (ΔH)
  • Heat energy is enthalpy
  • Change in heat (delta means change/displacement)
  • Change in potential energy from before and after the reaction
• Potential energy is found in the chemical bonds of molecules (opposite nucleus attracted to electron - holding onto each other to form a strong bond)
  • Intramolecular bonds (bonds within the molecule)
  • Attractive electrostatic force (- force)
  • • force - like charges - repels

COULOMB’S LAW

Coulomb’s Law Equation: $F = \frac{kQ_{1}Q_{2}}{r^{2}}$

• Q1 and Q2 refer to the charges of the particles
• Radius (denominator) is the distance apart (force approaches 0 as distance increases)
  • As distance decreases, force increases and bond forms between atoms
• Coulomb’s law (energy/heat interaction): $E = \frac{kQ_{1}Q_{2}}{r}$ since $F = \frac{E}{r}$
  • As r approaches infinity, energy approaches 0
  • As r decreases, energy becomes more negative - exothermic (if opposite charges, releases more energy) or more positive - endothermic (if like charges, absorbs more energy)

POTENTIAL ENERGY DIAGRAMS

• Bonds forming (coming together) is an exothermic (exo - out) process (energy is released to form the bond)

• Bonds breaking is an endothermic (endo - in) interaction (energy is added to separate the bonds)

  • Heating for decomposition is adding energy to break the bond

• Chemical reactions: the rearrangement of atoms

  • First bonds need to be broken (endothermic)
    • Ability to break bonds depends on atom (different atoms have different properties - some bonds are stronger/weaker, but all take energy to break and are equal in both directions - amount of energy to break = amount of energy to form)
  • New bonds can be formed (exothermic)
  • Process is shown in a potential energy diagram (PED)
    • Shows the enthalpy of the reaction (change in potential energy from before and after the reaction)
      • Actually is the net amount of heat change in the reaction from the start to the end (energy was added and energy was released in rearranging bonds - 2 steps)
        • Compares which part of the rearrangement had a greater impact (what step transfers most energy)
        • If the reaction releases more energy (negative enthalpy) than the amount need to break the bonds, exothermic (less potential energy at the end of the reaction)
        • If the reaction absorbs more energy (positive enthalpy) than the amount needed to break the bonds, endothermic (more potential energy at the end of the reaction)
      • Energy is released into or absorbed from surroundings - Bunsen burner, air in room, etc. (reaction is the system, some reactions need large surrounding energy to start reaction)
  • Spontaneous decompositions need very little energy to break the bond (PED has only a small increase in energy to break bond before new bonds are formed), so ambient energy starts the reaction