5. The Periodic Table
Reading time: ~32 minutes | Pages: 8
Quantum mechanics reveals the deep structure underlying chemistry's organization.
Quantum Basis of Periodicity
Periodic properties arise from electron shell structure:
- Pauli exclusion: Limits electrons per orbital
- Shell filling: Electrons occupy lowest available energy states
- Valence electrons: Outermost electrons determine chemistry
- Screening: Inner electrons shield outer electrons from nucleus
Shell Structure
| Shell (n) | Subshells | Max Electrons | Period |
|---|---|---|---|
| K (n=1) | 1s | 2 | 1 |
| L (n=2) | 2s, 2p | 8 | 2 |
| M (n=3) | 3s, 3p, 3d | 18 | 3, 4 |
| N (n=4) | 4s, 4p, 4d, 4f | 32 | 4, 5, 6 |
Maximum electrons per shell: $2n^2$
Groups and Families
Group 1 (Alkali metals): ns¹
- Li, Na, K, Rb, Cs, Fr
- One valence electron → highly reactive
- Easily lose electron to form +1 ions
Group 2 (Alkaline earth): ns²
- Be, Mg, Ca, Sr, Ba, Ra
- Two valence electrons → reactive
- Form +2 ions
Groups 3-12 (Transition metals): (n-1)d¹⁻¹⁰ ns¹⁻²
- Filling d orbitals
- Variable oxidation states
- Colored compounds, catalytic activity
Group 17 (Halogens): ns² np⁵
- F, Cl, Br, I, At
- One electron short of filled shell → very reactive
- Form -1 ions
Group 18 (Noble gases): ns² np⁶
- He, Ne, Ar, Kr, Xe, Rn
- Filled outer shell → chemically inert
Ionization Energy
Energy required to remove electron:
Trends:
- Across period (left to right): Increases (stronger nuclear attraction, less screening)
- Down group: Decreases (larger atoms, more screening)
- Peaks: Noble gases (filled shells)
- Valleys: Alkali metals (single valence electron)
Electron Affinity
Energy released when adding electron:
Trends:
- Halogens: Highest EA (one electron from filled shell)
- Noble gases: Negative EA (don't want extra electron)
- Generally increases: Left to right across period
Atomic Radius
Trends:
- Down group: Increases (more shells)
- Across period: Decreases (stronger nuclear pull with same shells)
Effective size estimate:
Electronegativity
Tendency to attract electrons in bond (Pauling scale):
- Fluorine: 4.0 (most electronegative)
- Cesium/Francium: ~0.7 (least electronegative)
- Trend: Increases diagonally toward top-right
Determines bond type:
- $\Delta EN > 1.7$: Ionic
- $0.4 < \Delta EN < 1.7$: Polar covalent
- $\Delta EN < 0.4$: Nonpolar covalent
Lanthanides and Actinides
Lanthanides (elements 57-71):
- Filling 4f orbitals
- Configuration: [Xe] 4f¹⁻¹⁴ 5d⁰⁻¹ 6s²
- Similar chemistry (hard to separate)
- Lanthanide contraction: radius decreases across series
Actinides (elements 89-103):
- Filling 5f orbitals
- Configuration: [Rn] 5f¹⁻¹⁴ 6d⁰⁻¹ 7s²
- Many are radioactive
- More variable oxidation states than lanthanides
Chemical Bonding Preview
Ionic bonding:
- Electron transfer: metal + nonmetal
- Example: Na⁺Cl⁻
- Driven by: Low IE (metal) + high EA (nonmetal)
Covalent bonding:
- Electron sharing: nonmetal + nonmetal
- Example: H₂, O₂, CH₄
- Quantum: Overlapping orbitals, bonding/antibonding states
Metallic bonding:
- Delocalized electrons (electron sea)
- Metal + metal
- Conductivity, malleability
Exceptions to Simple Patterns
- Chromium: [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²) - half-filled d stability
- Copper: [Ar] 3d¹⁰ 4s¹ (not 3d⁹ 4s²) - filled d stability
- Palladium: [Kr] 4d¹⁰ (no 5s electrons!)
- Helium: Group 18 but 1s² (not p⁶)
- Hydrogen: Unique - can be +1, -1, or share
Modern Periodic Law
Properties of elements are periodic functions of their atomic number (not atomic mass).
Quantum explanation:
- Atomic number = number of protons = number of electrons
- Electron configuration determines chemistry
- Pauli exclusion → shell structure → periodicity
Mendeleev (1869) organized by mass; Moseley (1913) proved atomic number is fundamental