2.2.1 Electron structure

shells

• electrons arranged around nucleus in principle quantum energy levels

• principal quantum numbers (n) used to number energy levels

  • lower number=closer to nucleus

how many electrons can each principle quantum energy level hold

• n=1→ 2

• n=2→ 8

• n=3→ 18

• n=4→32

subshells

• principal quantum shells split into subshells:

  • s holds 2 electrons

  • p holds 6 electrons

  • d holds 10 electrons

  • f (for elements with 57+ electrons) holds 14 electrons

orbitals

• subshells contain atomic orbitals

• orbitals exist at specific energy levels- electrons only found at these levels, not in-between

• each orbital can only hold 2 electrons

s orbital shape

• spherical shape

• size increases with shell number

p orbital shape

• dumbbell shape

• every shell has 3 p orbitals except 1st

• p-orbitals occupy x,y,z axes- perpendicular

• lobes are larger and longer with increasing shell number

filling orbitals

• electrons can spin either clockwise or anticlockwise

• electrons with same spin repel each other- spin-pair repulsion:

  • electrons will occupy separate orbitals in same subshell to minimise repulsion

  • electrons in an orbital must have opposite spin

ground state

• most stable electron configuration of an atom which has the lowest amount of energy

• 4s orbital filled and emptied before 3d orbital because it has lower energy

box diagrams

• each orbital can hold up to 2 electrons

• orbitals are modelled as boxes

• an electron is shown as a single headed arrow with either an up or down spin

chromium and copper

• elements are more stable if the d shell is partially or fully filled

• chromium and copper are one electron off having a fully or partially filled shell, so an electron is promoted to the d shell to complete it

  i.e. Cu → [Ar] 4s¹3d¹⁰ Cr → [Ar] 4s¹3d⁵

electron configuration of atoms

• block in which element is in tells us the which orbital the valence electrons are in

• period tells us which shell they are in.