Periodicity & Classification of Elements

The modern periodic table is based on the Modern Periodic Law (Moseley, 1913): physical and chemical properties of elements are periodic functions of their atomic number (not atomic weight). The table has 18 groups and 7 periods, organized by electronic configuration into four blocks (s, p, d, f).

Electronic Configuration and Block Classification

s-Block (Groups 1-2): $ns^{1}$-2. Metals. Low IE, electropositive, strong reducing agents. Include H and He (anomalous placements).

p-Block (Groups 13-18): ns^{2}$$np^{1}-6. Contains metals, metalloids, and non-metals. Most diverse chemistry. Group 18 = noble gases (fully filled valence shell).

d-Block (Groups 3-12): (n-1)$d^{1}$-10 $ns^{0}$-2. Transition metals (Groups 3-11 strictly, where incomplete d-subshell exists in element or common ion). Variable oxidation states, coloured compounds, catalytic activity, complex formation.

f-Block (Lanthanoids and Actinoids): (n-2)$f^{1}$-14. Inner transition metals. Lanthanoid contraction. Similar properties within each series.

Periodic Trends

Atomic Radius: Decreases across a period (increasing Zeff, electrons pulled closer). Increases down a group (new shell added). Covalent radius < van der Waals radius. For d-block: radius decreases slowly across the series, then slightly increases at the end (electron-electron repulsion in d-orbitals). Lanthanoid contraction causes post-lanthanoid elements (Hf, Ta, etc.) to have similar radii to their 4d counterparts (Zr, Nb, etc.).

Ionic Radius: Cations are smaller than parent atoms (fewer electrons, same nuclear charge). Anions are larger (more electrons, same nuclear charge). For isoelectronic species (same electron count), radius decreases with increasing nuclear charge: $O^{2}$- > F^- > Ne > Na⁺ > $Mg^{2}$+ > $Al^{3}$+.

Ionisation Energy (IE): Energy required to remove the outermost electron from a gaseous atom in ground state. IE₁ increases across a period (general trend with exceptions). IE₁ decreases down a group. Exceptions to period trend: (1) IE of B < Be (B loses 2p electron which is easier than losing 2s of Be — 2p is higher energy, further from nucleus). (2) IE of O < N (N has half-filled $2p^{3}$ — extra stability; O has paired electron in $2p^{4}$ which experiences repulsion, easier to remove).

Successive IE: IE₁ < IE₂ < IE₃ < ... (each subsequent removal is harder because ion has greater Zeff). A large jump in IE indicates the next electron is being removed from a core shell (used to identify group).

Electron Affinity (EA): Energy change when a gaseous atom gains an electron. Generally becomes more negative across a period and less negative down a group. Exceptions: EA of F < Cl (F is too small, high electron density causes repulsion for incoming electron). EA of N ≈ 0 (half-filled $2p^{3}$ resists additional electron). EA of noble gases ~ 0 (fully filled shell).

Electronegativity (EN): Tendency of an atom to attract shared electrons in a bond (Pauling scale). Increases across a period, decreases down a group. F is most electronegative (4.0). Cs is least electronegative among main group metals (0.7). EN determines bond polarity and is not defined for noble gases on the Pauling scale.

Metallic Character: Decreases across a period, increases down a group. Metals tend to lose electrons (low IE), non-metals tend to gain (high EA, high EN). The diagonal from B to At roughly separates metals from non-metals, with metalloids along the boundary.

Anomalous Properties of First Elements

The first element of each group (Li, Be, B, C, N, O, F) shows significant differences from the rest:

1. Small size → high charge density → predominantly covalent compounds
2. No d-orbitals → max covalence limited (4 for Period 2)
3. Strong ppi-ppi bonding → forms multiple bonds (C=C, C=O, N=N, N triple bond N) that heavier elements cannot
4. Diagonal relationship instead of vertical: Li-Mg, Be-Al, B-Si, C-P (similar charge/radius ratio)
5. High IE and EN compared to group members

The key problem-solving concept is identifying the correct periodic trend and its underlying electronic cause — particularly the exceptions (B < Be, O < N for IE; F < Cl for EA; isoelectronic series radius).

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