2.2.1

• Aufbau principle: electrons will occupy the lowest energy level available
• Orbital fill order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
• Pauli’s Exclusion Principle: up to 2 electrons in opposite spin can fill a sub shell
• Hund’s rule: every orbital will be singularly occupied before doubly occupied (as electrons repel each other)
• Exceptions to the orbital fill order: chromium and copper
• Chromium: [Ar] 4s^1 3d^5
• Chromium and Copper fill differently as they are more stable with half-full or full sub shells
• Copper: [Ar]4s^1 3d^10
• number of electrons in a shell

  2n^2
• s orbitals

  spherical, contain 2 electrons
• p orbitals

  dumb-bell shaped, contain 6 electrons
• d orbital

  contain 10 electrons
• metal subshell

  s
• non-metal subshell

  p
• transition metal subshell

  d
• atomic orbital

  region around the nucleus that can hold up to 2 electrons in opposite spin
• once the 3d subshell is full it falls below the 4s subshell