expectation that elements with those atomic masses would be discovered. Those blanks correspond to the elements we now know as scandium, gallium, germanium, and technetium. Figure: Mendeleev’s Periodic Table, as Published in the German Journal Annalen der Chemie und Pharmacie in 1872: The column headings “Reihen” and “Gruppe” are German for “row” and “group.” Formulas indicate the type of compounds formed by each group, with “R” standing for “any element” and superscripts used where we now use subscripts. Atomic masses are shown after equal signs and increase across each row from left to right. The most convincing evidence in support of Mendeleev’s arrangement of the elements was the discovery of two previously unknown elements whose properties closely corresponded with his predictions (Table below). Two of the blanks Mendeleev had left in his original table were below aluminum and silicon, awaiting the discovery of two as-yet-unknown elements, eka-aluminum and eka-silicon (from the Sanskrit eka, meaning “one,” as in “one beyond aluminum”). The observed properties of gallium and germanium matched those of eka-aluminum and eka-silicon so well that once they were discovered, Mendeleev’s periodic table rapidly gained acceptance. Property (predicted) Gallium (observed) eka-Silicon (predicted) Germanium (observed) atomic mass 68 69.723 72 72.64 element metal metal dirty-gray metal gray-white metal low mp* mp = 29.8°C high mp mp = 938°C d = 5.9 g/cm3 d = 5.91 g/cm3 d = 5.5 g/cm3 d = 5.323 g/cm3 oxide E2O3 Ga2O3 EO2 GeO2 Aluminum
is the most important for the lanthanidesAny of the 14 elements between Z=58 (cerium) and Z=71 (lutetium). and for most of the actinidesAny of the 14 elements between Z=90 (thorium) and Z=103 (lawrencium).. Summary The chemical families consist of elements that have the same valence electron configuration and tend to have similar chemistry. The alkali metals (group 1) have ns1 valence electron configurations and form M+ ions, while the alkaline earth metals (group 2) have ns2 valence electron configurations and form M2+ ions. Group 13 elements have ns2np1 valence electron configurations and have an overwhelming tendency to form compounds in the +3 oxidation state. Elements in group 14 have ns2np2 valence electron configurations but exhibit a variety of chemical behaviors because they range from a nonmetal (carbon) to metals (tin/lead). Carbon, the basis of organic compounds, has at least four allotropes with distinct structures: diamond, graphite, fullerenes, and carbon nanotubes. The pnicogens (group 15) all have ns2np3 valence electron configurations; they form compounds in oxidation states ranging from −3 to +5. The chalcogens (group 16) have ns2np4 valence electron configurations and react chemically by either gaining two electrons or by formally losing four or six electrons. The halogens (group 17) all have ns2np5 valence electron configurations and are diatomic molecules that tend to react chemically by accepting a single electron. The noble gases (group 18) are monatomic gases that are chemically quite unreactive due to the presence of a filled shell of electrons. The transition metals (groups 3–10) contain partially filled sets of d orbitals, and the lanthanides and the actinides are those groups in which f orbitals are being filled. These groups exhibit strong horizontal similarities in behavior. Many of the transition metals form M2+ ions, whereas the chemistry of the lanthanides and actinides is dominated by M3+ ions. Key Takeaway Periodic properties and the chemical behavior of the elements correlate strongly with valence electron configurations and Mulliken electronegativities. 7.5 Elements in Biological Systems