BTEC Applied Science Unit One: Chemistry - Electron Configuration

Overview

• Electron configuration describes the arrangement of electrons in an atom.
• Electrons orbit the nucleus in defined shells (n = 1, 2, 3, etc.).
• Each shell contains subshells noted as s, p, d, and f.
• Subshells are composed of orbitals, each of which can hold a maximum of two electrons.

Shells and Subshells

• First Three Shells
  • n = 1: Contains only s subshell.
  • n = 2: Contains s and p subshells.
  • n = 3: Contains s, p, and d subshells.
• Orbitals and Electron Capacity
  • s subshell: 1 orbital.
  • p subshell: 3 orbitals.
  • d subshell: 5 orbitals.
  • Each orbital can hold 2 electrons.

Electron Capacity by Shell

• First Shell (n=1):
  • Maximum 2 electrons (1s^2).
• Second Shell (n=2):
  • Maximum 8 electrons (2s^2 2p^6).
• Third Shell (n=3):
  • Maximum 18 electrons (3s^2 3p^6 3d^10).

Examples

• Fluorine (9 electrons):

  • Electron configuration: 1s² 2s² 2p⁵.
  • Fill lowest energy orbitals first, e.g., n=1 then n=2.

• Calcium (20 electrons):

  • Electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² (4s fills before 3d).

• Magnesium (Mg²⁺):

  • Ion with 10 electrons has the configuration: 1s² 2s² 2p⁶, same as Neon.

Tips and Tricks

• Electrons fill orbitals one at a time (up, up, up, then down, down, down) in p, d subshells.
• Visualize electron filling like people choosing seats on a bus - prefer empty pairs.

Practice

• Draw electron configuration diagrams for the first 20 elements and their ions.
• Understand group trends:
  • Group 1 elements form +1 ions.
  • Group 2 elements form +2 ions.
  • Group 7 elements form -2 ions.
  • Group 0 (noble gases) generally do not form ions.

Conclusion

• Understanding electron configuration is crucial for exams and practical understanding of chemistry.
• Practice by drawing diagrams and memorizing configurations for better retention.