6. Electronic Structure & Periodic Trends
Electrons live in orbitals described by four quantum numbers. How they fill — and how the resulting arrangements repeat across the periodic table — explains nearly all of chemical behaviour.
Quantum Numbers and Orbitals
- $n$ — principal: energy level (1, 2, 3, …)
- $l$ — angular momentum: subshell shape (0 = s, 1 = p, 2 = d, 3 = f)
- $m_l$ — magnetic: orbital orientation
- $m_s$ — spin: $+\tfrac{1}{2}$ or $-\tfrac{1}{2}$
Each orbital holds at most two electrons (Pauli exclusion), and degenerate orbitals fill singly first (Hund's rule).
Electron Configurations
Example: Iron, $Z = 26$: $$1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6 \quad \text{or} \quad [\text{Ar}]\,4s^2 3d^6$$
The 4s fills before 3d, but when iron forms Fe2+, the 4s electrons leave first: Fe2+ is $[\text{Ar}]\,3d^6$.
Periodic Trends
- Atomic radius: decreases across a period (rising effective nuclear charge), increases down a group (new shells).
- Ionization energy: generally increases across a period, decreases down a group. Exceptions: Be > B (full 2s subshell) and N > O (half-filled 2p subshell).
- Electronegativity: increases across a period, decreases down a group; fluorine is the most electronegative element.
Practice Problems
- Write the electron configuration of Cl−.
- Which has the larger atomic radius: Na or Mg? Explain.
- Rank Li, Be, and B by first ionization energy, lowest to highest.
- Which element is the most electronegative?
Answers: 1. [Ne] 3s2 3p6 2. Na — same shell, but Mg has more protons pulling the electrons in 3. Li < B < Be (Be's full 2s subshell makes it harder to ionize than B) 4. Fluorine