Classification of Elements and Periodicity in Properties - Periodic Trends in Properties of Elements

Periodic Trends in Atomic Radius: Across a period, the atomic radius decreases due to an increase in the effective nuclear charge ($Z_{eff}$) which pulls the electrons closer. Down a group, the atomic radius increases because new principal energy shells ($n$) are added.

Ionic Radius: Cations are smaller than their parent atoms (e.g., $Na^+ < Na$) due to increased $Z_{eff}$. Anions are larger than their parent atoms (e.g., $Cl^- > Cl$) because of increased inter-electronic repulsion and decreased $Z_{eff}$ per electron.

Isoelectronic Species: For atoms/ions with the same number of electrons, the radius decreases as the atomic number ($Z$) increases. Example: $O^{2-} > F^- > Na^+ > Mg^{2+}$.

Ionization Enthalpy (${\Delta}_i H$): It is the energy required to remove an electron from an isolated gaseous atom. It increases across a period (increasing $Z_{eff}$) and decreases down a group (increasing atomic size and shielding effect).

Electron Gain Enthalpy (${\Delta}_{eg} H$): The enthalpy change when an electron is added to a neutral gaseous atom. Halogens have highly negative ${\Delta}_{eg} H$. $Cl$ has a more negative ${\Delta}_{eg} H$ than $F$ due to less inter-electronic repulsion in $Cl$'s larger $3p$ orbital.

Electronegativity: The ability of an atom in a chemical compound to attract shared electrons. According to the Pauling scale, Fluorine ($F$) is the most electronegative ($4.0$) and Cesium ($Cs$) is the least.

Valency and Oxidation States: Valency across a period increases from $1$ to $4$ and then decreases to $0$ for noble gases. Elements in the same group usually exhibit the same valence.

Anomalous Properties of Second Period Elements: The first element of each group ($Li, Be, B, C, N, O, F$) shows different behavior due to small size, high electronegativity, high ionization enthalpy, and absence of vacant $d$-orbitals.