Part 1: Historical Development and Modern Periodic Law

The modern periodic table's intellectual history spans from Döbereiner's triads (1817) through Newlands' octaves (1865), Mendeleev's atomic-mass-based table (1869), and finally Moseley's 1913 revision. Moseley showed through X-ray spectroscopy that atomic number (not mass) is the fundamental property governing chemical behavior. The modern periodic law states: physical and chemical properties of elements are periodic functions of their atomic numbers. This resolved Mendeleev's anomalies (e.g., Ar/K and Co/Ni) and placed elements in their correct positions regardless of atomic mass.

Part 2: Structure of the Modern Periodic Table

The table has 7 periods and 18 groups. Period number = highest principal quantum number present. Group number for s- and p-block elements = total number of valence electrons. The four blocks are defined by the last-filled subshell:

• s-block (Groups 1–2): alkali and alkaline earth metals; highly reactive, low IE
• p-block (Groups 13–18): includes non-metals, metalloids, and noble gases
• d-block (Groups 3–12): transition metals; variable oxidation states
• f-block (below main table): inner transition metals; lanthanoid and actinoid series

Part 3: Atomic and Ionic Radii

Atomic radius decreases across a period (Zeff increase at constant n) and increases down a group (new shell). Cations are smaller than parent atoms; anions are larger. Isoelectronic species have radii determined solely by nuclear charge — more protons = smaller radius. The 10-electron isoelectronic series ($N^{3-}$ through $Al^{3+}$) is the standard NEET question type for this concept.

Part 4: Ionization Enthalpy

IE is endothermic and generally increases across a period and decreases down a group. Two Period-2 exceptions are essential for NEET: IE(Be) > IE(B) because Be's filled 2$s^{2}$ resists electron removal more than B's lone 2$p^{1}$; and IE(N) > IE(O) because N's half-filled 2$p^{3}$ is more stable than O's 2$p^{4}$ with its paired electron in one orbital. These exceptions arise from subshell stability overriding the Zeff effect.

Part 5: Electron Gain Enthalpy

EGE measures energy released (mostly exothermic for non-metals) when an electron is added to a gaseous atom. The paramount NEET exception: EGE(Cl) = −349 kJ/mol is more negative than EGE(F) = −328 kJ/mol, because F's compact 2p orbital causes excessive electron–electron repulsion. Noble gases have positive EGE; Groups 2 and 15 also deviate from the trend.

Part 6: Electronegativity and Diagonal Relationships

EN (Pauling scale) peaks at F (4.0). It increases across a period and decreases down a group — no major exception. EN difference determines bond type. Diagonal relationships (Li–Mg, Be–Al, B–Si) arise from comparable polarizing powers (charge-to-size ratio), causing similar chemical behaviors: both Be and Al oxides are amphoteric; both Li and Mg form normal (not super)oxides; both B and Si chlorides are hydrolyzed by water.