The significance of the periodic law

In 1869 D.I. Mendeleev stated the periodic law: the properties of elements depend periodically on their atomic mass (in the modern statement, on the nuclear charge, i.e. the atomic number). Then 63 elements were known, now 118; for undiscovered elements (gallium, scandium, germanium) Mendeleev left gaps in the table and predicted their properties. Element properties change regularly in the system: down a main group, atomic radius, metallic character and reducing power increase and electronegativity decreases; across a period from left to right, atomic radius decreases, metallic character and reducing power weaken, and non-metallic character, oxidising power and electronegativity increase. In a main group the number of outer electrons equals the group number, so the highest oxidation state equals the group number; the lowest oxidation state of a non-metal is group number − 8 (for example sulfur: +6 and −2). Across a period oxides change from basic through amphoteric to acidic, and hydroxides from bases to acids. Using these regularities, an element's position (period, group) is found from its electron configuration, and the formulas and properties of its compounds can be predicted.