🔬 By the end of this lesson, you’ll be able to:

• Describe energy levels, sublevels, and orbitals:
  • Energy levels (shells) are the main layers where electrons reside
  • Sublevels (s, p, d, f) divide energy levels into regions of varying energy
  • Orbitals are 3D spaces within sublevels where electrons are most likely found
  • Understand how each orbital holds a maximum of 2 electrons
• Apply key principles to determine electron arrangements:
  • Aufbau Principle – electrons fill the lowest energy orbitals first
  • Hund’s Rule – electrons occupy orbitals singly before pairing
  • Pauli Exclusion Principle – no two electrons in an atom can have the same set of quantum numbers
  • Use these rules to build accurate electron configurations for elements
• Write full and shorthand electron configurations:
  • Full configuration example: Oxygen → 1s² 2s² 2p⁴
  • Shorthand configuration using noble gases: Oxygen → [He] 2s² 2p⁴
  • Learn how to use the periodic table as a guide for electron filling order
• Explain periodic trends based on electron configurations:
  • Understand trends in atomic radius, ionization energy, and electronegativity
  • See how electron shells and sublevel filling influence element properties across periods and groups
• Predict bonding behavior based on valence electrons:
  • Valence electrons (outermost electrons) determine how atoms bond
  • Learn how elements form ionic, covalent, or metallic bonds based on electron sharing or transfer
  • Explore examples like sodium donating an electron to chlorine, forming NaCl