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Chemistry Chapter 5 The Periodic Law Mendeleev’s Periodic Table Dmitri Mendeleev Mendeleev – organized periodic table • Vertical columns in atomic mass order • Made some exceptions to place elements in rows with similar properties (Te and I) • Horizontal rows have similar chemical properties • Gaps for “yet to be discovered” elements • Left questions: why didn’t some elements fit in order of increasing mass? Why did some elements exhibit periodic behavior? Moseley • Discovered that periodic table was in atomic number order, not atomic mass order • Explained the Te-I anomaly Periodic Law • Physical and chemical properties of the elements are periodic functions of their atomic numbers Modern Periodic Table • Discovery of noble gases yields new family (Group 18 – aka inert gases) • Lanthanides (#58 - #71) • Actinides (#90 – #103) Periods and Blocks of the Periodic Table • Periods – horizontal rows • Groups/Families – vertical columns; these elements share similar chemical properties (they have the same number of valence electrons) • Blocks – periodic table can be broken into blocks corresponding to s, p, d, f sublevels Orbital filling table Blocks and Groups – s block • Group1 – “The alkali metals” • One s electron in outer shell • Soft, silvery metals of low density and low melting points • Highly reactive, never found pure in nature Blocks and Groups – s block • Group 2 – “Alkaline Earth Metals” • 2 s electrons in outer shell • Denser, harder, stronger, less reactive than Group 1 • Too reactive to be found pure in nature Periodic Table with Group Names The Properties of a Group: the Alkali Metals Easily lose valence electron (Reducing agents) React violently with water React with halogens to form salts Blocks and Groups – d block • • • • Groups 3 -12 Metals with typical metallic properties Referred to as transition metals Group number = sum of outermost s and d electrons Properties of Metals Metals are good conductors of heat and electricity Metals are malleable Metals are ductile Metals have high tensile strength Metals have luster Examples of Metals Potassium, K reacts with water and must be stored in kerosene Copper, Cu, is a relatively soft metal, and a very good electrical conductor. Zinc, Zn, is more stable than potassium Mercury, Hg, is the only metal that exists as a liquid at room temperature Blocks and Groups – p block • Groups 13-18 • Properties vary greatly – metals, metalloids, and nonmetals • Group 17 – halogens are most reactive of non metals • Group 18 – noble gases are NOT reactive Properties of Nonmetals Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element. Nonmetals are poor conductors of heat and electricity Nonmetals tend to be brittle Many nonmetals are gases at room temperature Examples of Nonmetals Sulfur, S, was once known as “brimstone” Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure Microspheres of phosphorus, P, a reactive nonmetal Properties of Metalloids Metalloids straddle the border between metals and nonmetals on the periodic table. They have properties of both metals and nonmetals. Metalloids are more brittle than metals, less brittle than most nonmetallic solids Metalloids are semiconductors of electricity Some metalloids possess metallic luster Silicon, Si – A Metalloid Silicon has metallic luster Silicon is brittle like a nonmetal Silicon is a semiconductor of electricity Other metalloids include: Boron, B Germanium, Ge Arsenic, As Antimony, Sb Tellurium, Te Blocks and Groups – f block • Lanthanides – shiny metals similar to group 2 • Actindes – all are radioactive; plutonium – lawrencium are man-made Determination of Atomic Radius: Half of the distance between nucli in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels Table of Atomic Radii Ionization Energy - the energy required to remove an electron from an atom Increases for successive electrons taken from the same atom Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group Outer electrons are farther from the nucleus Ionization of Magnesium Mg + 738 kJ Mg+ + eMg+ + 1451 kJ Mg2+ + eMg2+ + 7733 kJ Mg3+ + e- Table of 1st Ionization Energies Another Way to Look at Ionization Energy Ionic Radii Cations Anions Positively charged ions Smaller than the corresponding atom Negatively charged ions Larger than the corresponding atom Table of Ion Sizes Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons Electronegativities tend to increase across a period * more nuclear charge, more power to attract electrons Electronegativities tend to decrease down a group or remain the same * additional energy levels result in less attraction to the nucleus Periodic Table of Electronegativities