Chapter 10: Periodicity

The tenth chapter of Cambridge International AS and A Level Chemistry provides a comprehensive look into the fundamental concept of periodicity, focusing primarily on the observable trends and recurring patterns exhibited by the elements within Period Three, stretching from sodium through to argon. The discussion establishes the historical context of the Periodic Table, highlighting Dmitri Mendeleev’s initial organization based on atomic mass and the subsequent refinement to ordering elements by atomic number. Significant detail is dedicated to the periodic variation in physical properties. The atomic radius is shown to decrease across Period Three, an effect attributed to the consistently increasing positive nuclear charge attracting the outermost electrons more strongly while the internal shielding remains largely constant. Trends in ionic radius are split between cations (positive ions), which decrease in size from sodium ion to silicon ion, and anions (negative ions), which decrease in size from phosphorus ion to chloride ion; both sets of ions are significantly impacted by the increasing nuclear charge across the period. The behavior of melting point and electrical conductivity is explained by the transition in structure and bonding across the row. Metallic elements (sodium, magnesium, and aluminum) display increasing conductivity and melting points due to the rise in the number of delocalized electrons and the higher positive charge of the ions, strengthening the metallic bond. Silicon, positioned centrally, registers the highest melting point, characteristic of its giant molecular covalent structure. Moving toward the right, the non-metallic elements possess simple molecular structures, resulting in very low melting points that require minimal energy to overcome the weak instantaneous dipole-induced dipole forces. Furthermore, the first ionization energy generally increases across the period, though minor decreases occur between magnesium and aluminum (Group Two to Group Thirteen) and between phosphorus and sulfur (Group Fifteen to Group Sixteen). The chemical properties of these elements are explored through their reactions with oxygen, chlorine, and water. Sodium and magnesium react with water to form strongly alkaline and weakly alkaline solutions, respectively. A critical periodic trend is seen in the acid-base nature of the oxides, which change from basic (sodium oxide and magnesium oxide, having ionic bonding) to amphoteric (aluminum oxide, reacting with both acids and bases) to increasingly acidic (covalently bonded oxides of silicon, phosphorus, and sulfur). The behavior of Period Three chlorides in water is similarly linked to bonding. Ionic chlorides (like sodium chloride) simply dissolve, producing near-neutral solutions, whereas the covalent chlorides (including aluminum chloride, silicon chloride, and phosphorus five chloride) undergo hydrolysis, reacting with water to release hydrogen chloride gas and resulting in highly acidic solutions.