Week 3

• Electronic configuration


  Arrangement of electrons in an atom into orbitals
• Ground state


  Lowest possible energy state for a set of electrons
• Excited state


  Higher energy than ground state
• Excited state


  Require energy to move to excited state
• Degenerate


  Set of orbitals which have the same energy
• Aufbau principle


  • Fill electrons from lowest possible energy level first
• Hund's rule


  • Electrons cannot be paired until each orbital in the set contains one electron
• Pauli principle


  • Cannot have two electrons with same set of quantum numbers
• Order to fill subshells

  • 1s
  • 2s
  • 2p
  • 3s
  • 3p
  • 3d
  • 4s
  • 4p
  • 4d
  • 5s
  • 5p
  • 5d
  • 5f
• Capacity of orbitals


  • 2 electrons in s orbitals
  • 6 electrons in p orbitals
  • 10 electrons in d orbitals
  • 14 electrons in f orbitals
• Effective nuclear charge (Zeff)

  Difference between attraction and repulsion, difference between core and valence electrons
• Atomic radius


  • Size of entire atom, increases from top to bottom, increases from right to left
• Ionisation energy


  Amount of energy required to remove 1 mol of electrons, increases as atomic radius decreases and electrons are closer to nucleus
• Electron affinity


  Amount of energy associated with the addition of 1 mol of electrons
• Electronegativity


  Ability to attract electrons towards itself
• Ion size


  • Cations have radius smaller than neutral atom, anions have radius larger than neutral atom