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Periodicity And the Periodic Table The Big Idea • Periodic trends in the properties of atoms allow us to predict physical and chemical properties Development of the Periodic Table • The periodic table evolved over time as scientists discovered more useful ways to compare and organize the elements The Periodic Table • Mendeleev (1869) – first periodic table • Based on increasing atomic mass and repeating properties of elements • Had spaces for “missing” elements he predicted Henry Moseley • Discovered in 1914 that elements’ properties more closely followed their atomic number • Modern Periodic Table based on this discovery Periodic Law • Physical and chemical properties of elements are periodic functions of atomic numbers Classification of the Elements • Elements are organized into different blocks in the periodic table according to their electron configurations Organizing the Elements by Electron Configuration • Electrons in the highest principal energy level are called valence electrons. • All group 1 elements have one valence electron. Organizing the Elements by Electron Configuration • The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found. • The number of valence electrons for elements in groups 13-18 is ten less than their group number. Organizing the Elements by Electron Configuration What is periodicity? • Properties of the elements change in a predictable way as you move through the periodic table • These properties include – – – – – Atomic radius Octet Rule Ionic radius Ionization energy Electronegativity How are the elements organized? • Periodic Table: a complete chart of all elements in the universe • Arranged according to physical and chemical properties • Each box on the table contains the atomic number, atomic mass, and chemical symbol How are the elements organized? • Groups: also known as families; are the columns • Have similar properties • Some have specific names: – – – – Family 1: Alkali Metals Family 2: Alkaline Earth Metals Families 3 – 12: Transition Metals Family 13: Boron Family How are the elements organized? – – – – – Family Family Family Family Family 14: 15: 16: 17: 18: Carbon Family Nitrogen Family Oxygen Family Halogen Family Noble Gases • Periods: the rows on the periodic table – Do not have similar properties The Modern Periodic Table Metals • To the left of the stair-step line • 88 elements • Tend to lose electrons • Most reactive in the s block • Includes alkali metals and alkaline earth metals Properties of Metals • • • • • • Shiny luster Good conductors of heat Good conductors of electricity Usually solids at room temp Malleable Ductile gold lead nickel copper Nonmetals • • • • On right side of stair-step line Tend to gain electrons Most reactive group is halogens Least reactive is Noble gases Properties of Nonmetals • • • • • Dull luster Poor conductors of heat Poor conductors of electricity Brittle Many are gases at room temp Carbon (graphite) bromine sulfur Metalloids • On either side of stair-step line • Have properties of metals and nonmetals • Includes all elements that touch the line except Al and Po • Many are used in transistors antimony germanium Alkali Metals • Group 1 (except H) • All have only 1 valence electron • Most reactive metals; never found in pure state in nature • Soft, shiny, have relatively low melting points Alkaline Earth Metals • Group 2 • All have 2 valence electrons • Are the second most reactive metals; never found naturally in pure state • Harder, denser, stronger than alkali metals • Have higher melting points than alkali metals Transition Metals • Groups 3 – 12 • All have 1 or 2 valence electrons (in s sublevels) • Do not fit into any other group or family • Have many irregularities in their electron configurations Boron Family • Group 13 • Have 3 valence electrons • Boron is a metalloid, while all of the others are metals Carbon Family • Group 14 • All have four valence electrons • Carbon is a nonmetal; Si and Ge are metalloids; Sn and Pb are metals Nitrogen Family • Group 15 • All have 5 valence electrons (in s and p sublevels) • N and P are nonmetals; As and Sb are metalloids; Bi is a metal Oxygen Family • Group 16 • All have 6 valence electrons (in s and p sublevels) • All are nonmetals except Te, which is a metalloid, and Po, which is a metal. Halogens • Means “salt former” • Group 17 • All have 7 valence electrons (in s and p sublevels) • Most reactive nonmetals • All are nonmetals except At, which is a metalloid Noble Gases • Group 18 • Complete, stable electron configuration (no valence electrons) • Most unreactive elements Rare Earth Elements • Found in 2 rows at bottom of Periodic Table • Also known as the inner transition metals • Lanthanide series: starts with La • Actinide series: starts with Ac • Little variation in properties • Actinides are radioactive; only first three and Pu are found in nature Periodic Trends • Trends among elements in the periodic table include their size and their ability to lose or attract electrons Atomic Radius • For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms. Atomic Radius Atomic Radius • Atomic radius generally increases as you move down a group. • The outermost orbital size increases down a group, making the atom larger. Ionic Radius • An ion is an atom or bonded group of atoms with a positive or negative charge. • When atoms lose electrons and form positively charged ions, they always become smaller for two reasons: 1. The loss of a valence electron can leave an empty outer orbital resulting in a small radius. 2. Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius. Ionic Radius • When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion. Ionic Radius Ionic Radius • The ionic radii of positive ions generally decrease from left to right. • The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16. Ionization Energy • The energy needed to remove one of its electrons • Decreases as you move down a group • Increases as you move across a period • Successive ionization energies increase for every electron removed Ionization Energy Octet Rule (Rule of 8) • Atoms tend to gain, share, or lose in order to acquire a full set of valence electrons (in most cases, this is 8) Electronegativity • Reflects an atom’s ability to attract electrons in a chemical bond • Related to its ionization energy and electron affinity • Increases as you move across a period • Increases as you move up a group Electronegativity