Topic 3 explores how physical and chemical properties vary systematically across periods and down groups, based on electron configuration and nuclear charge.
Arranged by increasing atomic number. Periods: horizontal rows (same number of electron shells). Groups: vertical columns (same number of valence electrons). Metals (left), non-metals (right), metalloids (staircase). s-block (Groups 1-2), p-block (13-18), d-block (transition metals, 3-12).
Atomic radius decreases across a period (increasing nuclear charge pulls electrons closer) and increases down a group (more shells). IE increases across a period and decreases down a group. Electronegativity (Pauling scale) increases across and up — fluorine is most electronegative (4.0).
Melting points: Na, Mg, Al increase (metallic bonding strengthens). Si is very high (giant covalent). P₄, S₈, Cl₂, Ar are low (simple molecular with weak IMF). Period 3 oxides: Na₂O, MgO are basic (ionic); Al₂O₃ is amphoteric; SiO₂ is weakly acidic; P₄O₁₀, SO₃, Cl₂O₇ are acidic.
d-block elements with incomplete d subshells (Sc to Zn, though Zn is debated). Properties: variable oxidation states (due to close 3d and 4s energies), coloured compounds (d-d transitions), catalytic activity, complex ion formation. Common oxidation states: Fe²⁺/³⁺, Cu⁺/²⁺, Mn²⁺/⁴⁺/⁷⁺, Cr³⁺/⁶⁺.
The periodic table organises all chemical knowledge. Understanding trends allows you to predict properties of unfamiliar elements: bonding type, oxide acidity/basicity, reactivity, and compound formulas. It connects Topics 2 (atomic structure), 3 (periodicity), 4 (bonding), and beyond. Exam questions frequently ask you to explain trends using nuclear charge, shielding, and electron configuration.
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