Unit 5: Electrons
Electrons don't orbit in fixed circles — they occupy regions of probability called ORBITALS. Each orbital can hold a maximum of 2 electrons (Pauli Exclusion Principle).
Shells: Numbered 1, 2, 3, etc. (the energy level). Higher = farther from nucleus = higher energy.
Subshells within each shell: s, p, d, f
• s subshell: 1 orbital → holds 2 electrons; spherical shape
• p subshell: 3 orbitals → holds 6 electrons; dumbbell-shaped along x, y, z axes
• d subshell: 5 orbitals → holds 10 electrons
Nodes: Points in an orbital where there is ZERO probability of finding an electron.
• 1st shell: 0 nodes
• 2nd shell: 1 node
• 3rd shell: 2 nodes
(Each shell has one more node than the previous)
Orbital filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p...
(Note: 4s fills BEFORE 3d)
An electron configuration describes exactly where all the electrons in an atom are located.
Full notation example — Sodium (Na, Z=11):
1s² 2s² 2p⁶ 3s¹
(Read: 1s subshell with 2 electrons, 2s with 2, 2p with 6, 3s with 1)
Total: 2+2+6+1 = 11 electrons ✓
Noble gas shorthand: Abbreviated by replacing inner electrons with the nearest noble gas in brackets.
Sodium: [Ne] 3s¹ (Neon has the configuration 1s²2s²2p⁶)
Rules:
1. Fill lowest energy orbitals first (Aufbau principle)
2. Maximum 2 electrons per orbital (Pauli Exclusion Principle)
3. Fill equal-energy orbitals singly before pairing (Hund's rule)
Examples:
• Hydrogen (Z=1): 1s¹
• Helium (Z=2): 1s²
• Carbon (Z=6): 1s² 2s² 2p²
• Oxygen (Z=8): 1s² 2s² 2p⁴
• Chlorine (Z=17): 1s² 2s² 2p⁶ 3s² 3p⁵ or [Ne] 3s² 3p⁵
Valence electrons: electrons in the OUTERMOST shell. These electrons determine chemical behavior.
• Elements in Group 1 have 1 valence electron
• Elements in Group 2 have 2 valence electrons
• Elements in Groups 13-18 have 3-8 valence electrons
Bohr Diagram: shows the nucleus with protons/neutrons labeled, and electrons arranged in concentric rings (shells).
• Shell 1: holds up to 2 electrons
• Shell 2: holds up to 8 electrons
• Shell 3: holds up to 8 electrons (for first 18 elements)
Useful for: showing which shell electrons are in, comparing potential energy and shielding.
Lewis Dot Structure: shows only the VALENCE ELECTRONS as dots around the element symbol.
• Dots are placed on the four sides (top, bottom, left, right) of the symbol
• First place one dot on each side before pairing up
Useful for: predicting bonding, because only valence electrons are shown (no distracting inner electrons).
Ions form when atoms GAIN or LOSE electrons.
Cation (positive ion): atom LOSES electrons → fewer electrons than protons → net + charge
• Metals typically form cations
• Example: Na (11 e⁻) → Na⁺ (10 e⁻), losing 1 electron
• Example: Mg (12 e⁻) → Mg²⁺ (10 e⁻), losing 2 electrons
Anion (negative ion): atom GAINS electrons → more electrons than protons → net − charge
• Nonmetals typically form anions
• Example: Cl (17 e⁻) → Cl⁻ (18 e⁻), gaining 1 electron
• Example: O (8 e⁻) → O²⁻ (10 e⁻), gaining 2 electrons
Predicting ion charge: elements gain/lose electrons to achieve a noble gas electron configuration (full outer shell — the octet rule).
Naming ions:
• Monatomic cations: element name + ion charge (if variable), e.g., Na⁺ = sodium ion; Fe²⁺ = iron(II) ion
• Monatomic anions: element name with -ide suffix, e.g., Cl⁻ = chloride; O²⁻ = oxide
Electron configurations and diagrams of ions:
• Draw the same as neutral atoms, but add/remove electrons from the outermost shell