Periodicity - The periodic table

The Periodic Table is arranged by increasing atomic number ($Z$). Horizontal rows are called periods and vertical columns are called groups.

Elements in the same group have the same number of valence electrons and share similar chemical properties. For example, Group 1 elements have a valence configuration of $ns^1$.

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

Atomic Radius: Decreases across a period (due to increasing $Z_{eff}$ pulling electrons closer) and increases down a group (due to the addition of new principal energy levels, $n$).

Ionic Radius: Cations ($M^{n+}$) are always smaller than their parent atoms because they lose electrons and sometimes an entire outer shell. Anions ($X^{n-}$) are always larger than their parent atoms due to increased electron-electron repulsion.

First Ionization Energy ($IE_1$): The energy required to remove one mole of electrons from one mole of gaseous atoms. It generally increases across a period and decreases down a group.

Electronegativity: A measure of the ability of an atom at attract a bonding pair of electrons in a covalent bond. It increases across a period and decreases down a group. Fluorine ($F$) is the most electronegative element.

Electron Affinity: The energy change when one mole of electrons is added to one mole of gaseous atoms. It generally becomes more exothermic (more negative) across a period.

Metallic Character: Decreases across a period and increases down a group. Metals generally have low $IE$ and low electronegativity, tending to form cations.

Oxide Periodicity: Across Period 3, oxides transition from basic ($Na_2O$, $MgO$), to amphoteric ($Al_2O_3$), to acidic ($SiO_2$, $P_4O_{10}$, $SO_3$, $Cl_2O_7$).