Periodic Trends Summary — Atomic Properties Across the Periodic Table
| Property | Definition | Unit | Trend Across Period (→) | Trend Down Group (↓) | Why | Highest / Largest | Lowest / Smallest |
|---|---|---|---|---|---|---|---|
| Atomic radius | Distance from nucleus to outermost electron shell | pm | Decreases | Increases | Across: increasing nuclear charge pulls electrons closer. Down: additional electron shells increase size. | Cs (265 pm) | He (31 pm) |
| Ionic radius (cations) | Radius of a positive ion | pm | Decreases (for same charge) | Increases | Cations are smaller than parent atoms due to lost electrons and increased effective nuclear charge. | Fr⁺ (194 pm) | Be²⁺ (27 pm) |
| Ionic radius (anions) | Radius of a negative ion | pm | Decreases | Increases | Anions are larger than parent atoms due to added electrons and increased electron-electron repulsion. | I⁻ (220 pm) | F⁻ (133 pm) |
| First ionization energy | Energy to remove one electron from gaseous atom | kJ/mol | Increases | Decreases | Across: greater nuclear charge holds electrons tighter. Down: outer electrons farther from nucleus, more shielded. | He (2372 kJ/mol) | Cs (376 kJ/mol) |
| Electron affinity | Energy change when atom gains an electron | kJ/mol | Becomes more negative (more exothermic) | Becomes less negative (less exothermic) | Across: higher nuclear charge attracts added electron more. Down: added electron is farther from nucleus. | Cl (–349 kJ/mol) | Noble gases (positive EA) |
| Electronegativity | Ability of bonded atom to attract shared electrons | Pauling scale | Increases | Decreases | Across: stronger nuclear charge attracts bonding electrons. Down: larger atoms hold bonding electrons less tightly. | F (3.98) | Cs (0.79) |
| Metallic character | Tendency to lose electrons and form cations | — | Decreases | Increases | Across: harder to lose electrons as nuclear charge increases. Down: easier to lose outer electrons farther from nucleus. | Fr / Cs | F |
| Nonmetallic character | Tendency to gain electrons and form anions | — | Increases | Decreases | Opposite of metallic character. Higher nuclear charge favors electron gain. | F | Fr / Cs |
| Melting point | Temperature at which solid becomes liquid | K | Irregular (increases then decreases) | Varies by group | Depends on bonding type: metals increase with bond strength, nonmetals vary. No simple universal trend. | W (3695 K) | He (0.95 K) |
| Electron shielding | Reduction of nuclear charge felt by outer electrons | — | Roughly constant (within same shell) | Increases | Down: more inner electron shells shield outer electrons from the nucleus more effectively. | Heavy elements (many shells) | H (no shielding) |
These trends apply cleanly to s-block and p-block (main-group) elements. The d-block and f-block deviate because d and f electrons shield the nucleus poorly, leaving Z_eff nearly flat across long stretches (transition-metal IEs barely move, the lanthanide contraction shrinks Hf to nearly Zr's size). Three named exceptions you'll be asked about: (1) IE dips at Groups 13 and 16 from subshell effects (B<Be, O<N); (2) electron affinity of F is less negative than Cl because F is so small the incoming electron sees strong e⁻–e⁻ repulsion; (3) the diagonal relationship — Li/Mg, Be/Al, B/Si pairs share chemistry because rightward-and-down moves cancel. Source: NIST ASD, CRC Handbook (97th ed.), Atkins' Physical Chemistry.
Frequently Asked Questions
Why do noble gases not follow the electronegativity trend?
Pauling defined electronegativity from bond dissociation energies, so it requires the element to actually form bonds. Helium, neon, and argon don't form stable compounds at all, so they have no Pauling value. Krypton, xenon, and radon do form compounds (XeF₂, XeF₄, KrF₂) and have been assigned Pauling values around 3.0, 2.6, and 2.2 — comparable to nitrogen or sulfur. Allred-Rochester and Mulliken scales, which compute electronegativity from atomic properties, do give numbers for all noble gases but are less commonly cited.
What is the diagonal relationship in the periodic table?
Three pairs in periods 2 and 3 share unexpected chemistry: Li/Mg, Be/Al, B/Si. Lithium burns in N₂ to form Li₃N like Mg₃N₂ (other Group 1 metals don't); beryllium hydroxide is amphoteric like Al(OH)₃; boron and silicon both form covalent network solids and weakly acidic oxides. The cause: moving right shrinks atoms and raises electronegativity, moving down enlarges them and lowers electronegativity, so a diagonal step largely cancels both effects, leaving the diagonal partner with a similar charge density and bonding character.
Why is effective nuclear charge important for understanding trends?
Z_eff is the net nuclear charge a valence electron actually feels after inner-shell electrons screen the proton count. It's the single concept underlying every trend in this table. Across a period, each added proton is screened only weakly (Slater: ~0.35 per same-shell electron), so Z_eff climbs and atoms shrink, IE rises, and EN rises. Down a group, each new shell adds nearly full unit screening from the inner core, so Z_eff stays roughly constant while atomic radius jumps from the new shell — IE and EN drop. Compute it with Slater's rules to see exact numbers.