The Modern Periodic Law and Table Structure

The classification of elements reached its modern form through Henry Moseley's landmark 1913 discovery that the frequencies of X-ray spectra of elements vary systematically with atomic number, not atomic mass. This gave birth to the Modern Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic numbers. This replaced Mendeleev's 1869 arrangement by atomic mass, which, while brilliant in predicting gaps for undiscovered elements, had inconsistencies (e.g., Co placed before Ni despite higher mass).

The long form of the periodic table contains 7 periods (horizontal rows) and 18 groups (vertical columns). Each period corresponds to the filling of a new principal quantum shell. The four blocks — s, p, d, and f — are defined by which subshell receives the differentiating electron:

Periodic Trends: The Zeff Framework

All major periodic trends are explained by two competing factors: effective nuclear charge ($Z_{eff} = Z - \sigma$) and the principal quantum number (n). Across a period, Z increases while n stays constant, so Zeff rises and electrons are pulled closer to the nucleus. Down a group, n increases (new shell added), which dominates — atoms become larger despite the rising Z.

Atomic radius decreases left to right across a period and increases top to bottom within a group. Ionic radii follow a related logic: cations are always smaller than their parent atoms (loss of electrons raises Zeff per remaining electron), and anions are always larger (extra electrons increase repulsion and expand the cloud). For isoelectronic species — ions with identical electron counts — radius is determined purely by nuclear charge: more protons on the same electron population = smaller radius. The benchmark 10-electron series is:

$\text{N}^{3-}\ (146\ \text{pm}) > \text{O}^{2-}\ (140\ \text{pm}) > \text{F}^{-}\ (133\ \text{pm}) > \text{Na}^{+}\ (98\ \text{pm}) > \text{Mg}^{2+}\ (66\ \text{pm}) > \text{Al}^{3+}\ (51\ \text{pm})$

Ionization Enthalpy: The Critical Exceptions

Ionization enthalpy (IE) is the energy required to remove the outermost electron from a gaseous atom. It generally increases across a period and decreases down a group. However, two essential exceptions occur in Period 2:

1. IE(Be) = 899 kJ/mol > IE(B) = 800 kJ/mol — Be loses an electron from its fully filled 2$s^{2}$ subshell, which has extra stability. B loses an electron from the singly occupied 2$p^{1}$ subshell, which is easier.
2. IE(N) = 1402 kJ/mol > IE(O) = 1314 kJ/mol — N has a half-filled 2$p^{3}$ configuration (one electron per orbital; minimum repulsion; extra stability by Hund's rule). O's 2$p^{4}$ forces one orbital to hold two electrons; the mutual repulsion makes that electron easier to remove.

Electron Gain Enthalpy: The Fluorine Paradox

Electron gain enthalpy (EGE) is the enthalpy change when a gaseous atom gains an electron. It generally becomes more negative across a period. The most important exception is:

EGE(Cl) = −349 kJ/mol is more negative than EGE(F) = −328 kJ/mol

Despite fluorine being the most electronegative element and having the highest Zeff among halogens, fluorine's 2p orbital is so compact that the incoming electron experiences severe electron–electron repulsion. Chlorine's 3p orbital is larger, so the electron is accommodated with less repulsion and more energy is released. This is the most frequently tested anomaly in NEET for this chapter. Noble gases have positive EGE; Groups 2 and 15 show less negative EGE due to filled and half-filled subshell stabilities respectively.

Electronegativity

Electronegativity (EN) on the Pauling scale measures the tendency of a bonded atom to attract shared electrons. Fluorine (EN = 4.0) is the most electronegative element. EN increases left to right across a period (Zeff rises, atomic size shrinks) and decreases down a group (larger atom, more shielding). The EN difference between bonded atoms determines whether a bond is nonpolar covalent (<0.5), polar covalent (0.5–1.7), or ionic (>1.7).

Diagonal Relationships

A unique pattern in the periodic table is the diagonal relationship: elements diagonally adjacent (one period lower, one group to the right) share strikingly similar properties because they have comparable charge-to-size (polarizing power) ratios. The three key pairs tested in NEET are:

NEET Focus Points

The first element in each group (H, Li, Be, B, C, N, O, F) often behaves anomalously because of its small size, high electronegativity, and absence of d-orbitals (cannot expand its octet beyond 4 bonds). Helium, despite having $1s^2$ like other Group 2 elements, is placed in Group 18 because it is chemically inert. For NEET, the three most tested exceptions — IE(Be) > IE(B), IE(N) > IE(O), and EGE(Cl) > EGE(F) — each have a common root cause: subshell stability (filled or half-filled) overrides the general Zeff-driven trend.