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Element Comparison & Data Check
Select two elements to compare the bundled teaching values and discuss broad periodic patterns.
Which element has the larger radius value in this teaching dataset?
Scientific Principles & Trend Fundamentals
Understanding broad periodic patterns, shielding, attraction, and important exceptions. Values in this page are a teaching dataset: atomic radius depends on the radius definition, heavy/superheavy values may be estimates, and the metallic-character percentage is an illustrative index rather than a standardized measured quantity.
Broad trend: Atomic size often decreases across a period and increases down a group, but radius definitions and local exceptions matter.
Why across a period? Protons are added to the nucleus while electrons fill the same valence shell. The increasing nuclear charge, only partly offset by shielding, tends to increase the attraction felt by valence electrons and pull the electron cloud inward.
Why down a group? Principal quantum number ($n$) increases, adding new electron shells farther from the nucleus.
Broad trend: Electronegativity usually rises across a period and falls down a group for comparable main-group elements; assigned values depend on the electronegativity scale and bonding context.
Why? Reflects an atom's ability to attract shared electrons in a bond. Smaller atoms with high $Z_{eff}$ (like Fluorine) attract bonding electrons strongly.
Important note: Electronegativity is a bonding concept. Introductory Pauling tables often leave lighter noble gases unassigned, while some sources/scales report values for heavier noble gases.
Broad trend: First ionization energy generally rises across a period and falls down a group, with well-known subshell and electron-pairing exceptions.
Anomalies: Not perfectly smooth!
• Group 2 vs Group 13 (e.g. Be to B): Beryllium's $2s^2$ full subshell is more stable than Boron's lone $2p^1$ electron, making Boron slightly easier to ionize.
• Group 15 vs Group 16 (e.g. N to O): Nitrogen has a half-filled $2p^3$ subshell. Oxygen's fourth $p$ electron experiences electron-electron pairing repulsion, slightly reducing its ionization energy.
Broad trend: Metallic behavior generally becomes less pronounced across a main-group period and more pronounced down many groups. This page’s 0–100 metallic index is an illustrative visualization, not a standardized physical scale.
Why? Metals readily lose valence electrons due to lower ionization energies and larger radii. Nonmetals tend to gain or share electrons due to high $Z_{eff}$ and electronegativity.