Periodicity - Periodic table

The Periodic Table is arranged by increasing atomic number ($Z$). Periods are horizontal rows (indicating the number of occupied main energy levels), and Groups are vertical columns (elements with the same number of valence electrons and similar chemical properties).

Elements are categorized into blocks: $s$, $p$, $d$, and $f$, representing the sub-shell being filled by the valence electrons.

Effective Nuclear Charge ($Z_{eff}$) is the net positive charge experienced by valence electrons. It increases across a period because the number of protons increases while inner-shell shielding remains constant. It remains approximately constant down a group.

Atomic Radius: Decreases across a period due to increased $Z_{eff}$ pulling electrons closer to the nucleus. Increases down a group as more main energy levels are added.

Ionic Radius: Cations ($M^+$) are smaller than their parent atoms because they lose electrons and often an entire outer energy level. Anions ($X^-$) are larger than their parent atoms due to increased electron-electron repulsion with the same nuclear charge.

First Ionization Energy ($IE_1$): The energy required to remove one mole of electrons from one mole of gaseous atoms: $X(g) \rightarrow X^+(g) + e^-$. It increases across a period and decreases down a group.

Electronegativity: The measure of an atom's ability to attract a shared pair of electrons in a covalent bond. Trends follow $IE_1$, increasing across a period and decreasing down a group.

Electron Affinity: The energy change when one mole of electrons is added to one mole of gaseous atoms: $X(g) + e^- \rightarrow X^-(g)$. Most elements have an exothermic first electron affinity.

Oxide Trends (Period 3): Oxides transition from basic ($Na_2O$, $MgO$) to amphoteric ($Al_2O_3$) to acidic ($SiO_2$, $P_4O_{10}$, $SO_2/SO_3$, $Cl_2O_7$).