Periodic Classification of Elements — Karnataka (SSLC) Class 10 Science Solutions (Free)

Periodic table = systematic arrangement of elements by atomic number in rows (periods) and columns (groups). Necessity: 1) Over 100 elements exist; classification helps organize and predict properties. 2) Elements in same group (vertical column) have similar chemical properties due to same valence electrons. 3) Helps predict reactivity, electronegativity, formation of compounds. Mendeleev left gaps for undiscovered elements, predicting their properties accurately.

Atomic number = number of protons = number of electrons (in neutral atom). Position: Group number = number of valence electrons (rightmost column with H). Period number = number of electron shells. Example: Na (Z=11): electron config 2,8,1; period 3 (3 shells), group 1 (1 valence electron). This arrangement explains why elements in same group react similarly; valence electrons determine bonding.

Metals: left side of periodic table, shiny, conduct electricity, form cations (Na, Fe, Cu, Al). Non-metals: right side, dull, poor conductors, form anions or share electrons (C, O, N, S, Cl). Metalloids: along diagonal between metals and non-metals, intermediate properties (B, Si, As, Sb). Metals are most abundant; non-metals essential for life (C, N, O); metalloids used in semiconductors (Si in computers).

Valence electrons = outermost electrons in electron shell; determine chemical behavior and bonding. Elements in same group (vertical column) have same number of valence electrons: Group 1 has 1 (alkali metals), Group 17 has 7 (halogens), Group 18 has 8 (noble gases, inert). Same valence electrons = similar reactivity and similar compound formation. Example: Li, Na, K all have 1 valence electron, all react with water forming hydroxide.

Periodic trends = properties change systematically across period or down a group. Electronegativity (tendency to attract electrons): increases left-to-right across period, decreases down a group (F highest at 4.0, Cs lowest at 0.79). Ionization energy (energy to remove electron): increases left-to-right, decreases down group. Reason: nuclear charge increases left-to-right, atomic radius increases down group. These trends help predict reactivity and bonding type.

Atomic radius = distance from nucleus to outermost electron. Across a period: radius DECREASES left-to-right because nuclear charge increases while electrons added to same shell (stronger attraction). Down a group: radius INCREASES because new electron shells added, increasing distance from nucleus (shielding effect of inner electrons). Example: Na > Mg > Al > Si (across period 3); Na < K < Rb < Cs (down group 1). Atomic radius relates to physical and chemical properties.

Sodium (Na, Z=11): electron config 2,8,1; period 3, group 1; valence electrons = 1. Chlorine (Cl, Z=17): electron config 2,8,7; period 3, group 17; valence electrons = 7. Oxygen (O, Z=8): electron config 2,6; period 2, group 16; valence electrons = 6. Na forms Na+1 (loses 1 electron), Cl forms Cl-1 (gains 1), O forms O-2 (gains 2). Period = shells, group = valence electrons.