Henry Moseley proposed the modern periodic law in 1913, stating that properties of elements are periodic functions of their atomic numbers, replacing Mendeleev's atomic-mass-based arrangement. The long-form periodic table has 7 periods and 18 groups, divided into s-, p-, d-, and f-blocks based on the subshell receiving the last electron. Atomic radius decreases across a period due to increasing effective nuclear charge and increases down a group as new electron shells are added. Cations are smaller than their parent atoms while anions are larger; for isoelectronic species, radius decreases as proton number increases. Ionization enthalpy generally increases across a period and decreases down a group, but two critical Period-2 exceptions exist: IE(Be) > IE(B) because Be loses from a stable filled 2$s^{2}$, and IE(N) > IE(O) because N's half-filled 2$p^{3}$ has extra stability. Electron gain enthalpy becomes more negative across a period, but chlorine (−349 kJ/mol) has a more negative EGE than fluorine (−328 kJ/mol) because F's compact 2p orbital causes severe electron–electron repulsion on addition of an extra electron. Fluorine is the most electronegative element with a Pauling scale value of 4.0, and electronegativity increases across periods and decreases down groups. Diagonal relationships between Li–Mg, Be–Al, and B–Si arise from comparable charge-to-size ratios, causing these pairs to share chemical properties despite being in different groups. Noble gases have positive EGE, and Groups 2 and 15 show deviations from the general EGE trend due to filled and half-filled subshell stability. The three most exam-critical exceptions — IE(Be)>IE(B), IE(N)>IE(O), and EGE(Cl)>EGE(F) — all share one root cause: subshell stability (filled or half-filled) overriding the otherwise dominant Zeff effect.