Chapter 3: Classification of Elements and Periodicity in PropertiesClass 11 Chemistry — summary, notes, extra questions & MCQ quiz

Summary

This unit traces the development of the periodic law and the modern periodic table. Early attempts — Döbereiner's triads, Newlands' law of octaves and de Chancourtois' telluric helix — preceded the independent work of Mendeleev and Lothar Meyer, who arranged elements by increasing atomic weight so that similar properties recurred periodically. Mendeleev left gaps and successfully predicted eka-aluminium (gallium) and eka-silicon (germanium). Moseley later showed that atomic number, not atomic weight, is the fundamental property, giving the modern periodic law: the properties of elements are periodic functions of their atomic numbers. The long-form table has seven periods and eighteen groups, and elements are classified by the orbital being filled into s-, p-, d- and f-blocks; the f-block lanthanoids and actinoids are inner-transition elements. Over three-quarters of the elements are metals, with metalloids (Si, Ge, As, Sb, Te) bordering the diagonal zig-zag line, and superheavy elements (\(Z>100\)) named by IUPAC numerical roots. Periodic trends arise from the interplay of nuclear charge and shielding: atomic and ionic radii decrease across a period and increase down a group, while ionisation enthalpy and electronegativity increase across and decrease down. Chemical reactivity is greatest at the two extremes of a period and least in the centre, and the nature of oxides changes from basic through amphoteric to acidic across a period.

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