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I. Development of the Modern Periodic Table (p. 174 - 181)

Ch. 6 - The Periodic Table & Periodic Law. I. Development of the Modern Periodic Table (p. 174 - 181). A. Mendeleev. Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass Elements with similar properties were grouped together There were some discrepancies.

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I. Development of the Modern Periodic Table (p. 174 - 181)

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  1. Ch. 6 - The Periodic Table & Periodic Law I. Development of the Modern Periodic Table(p. 174 - 181)

  2. A. Mendeleev • Dmitri Mendeleev (1869, Russian) • Organized elements by increasing atomic mass • Elements with similar properties were grouped together • There were some discrepancies

  3. A. Mendeleev • Dmitri Mendeleev (1869, Russian) • Predicted properties of undiscovered elements

  4. B. Moseley • Henry Moseley (1913, British) • Organized elements by increasing atomic number • Resolved discrepancies in Mendeleev’s arrangement • This is the way the periodic table is arranged today!

  5. C. Modern Periodic Table • Group (Family) • Period

  6. 1. Groups/Families • Vertical columns of periodic table • Numbered 1 to 18 from left to right • Each group contains elements with similar chemical properties

  7. 2. Periods • Horizontal rows of periodic table • Periods are numbered top to bottom from 1 to 7 • Elements in same period have similarities in energy levels, but not properties

  8. 3. Blocks • Main Group Elements • Transition Metals • Inner Transition Metals

  9. Lanthanides - part of period 6 Actinides - part of period 7 3. Blocks Overall Configuration

  10. Ch. 6 - The Periodic Table II. Classification of theElements(pages 182-186)

  11. A. Metallic Character • Metals • Nonmetals • Metalloids

  12. 1. Metals • Good conductors of heat and electricity • Found in Groups 1 & 2, middle of table in 3-12 and some on right side of table • Have luster, are ductile and malleable

  13. a. Alkali Metals • Group 1 • 1 Valence electron • Very reactive • Electron configuration • ns1 • Form 1+ ions • Cations • Examples: Li, Na, K

  14. b. Alkaline Earth Metals • Group 2 • Reactive (not as reactive as alkali metals) • Electron Configuration • ns2 • Form 2+ ions • Cations • Examples: Be, Mg, Ca, etc

  15. c. Transition Metals • Groups 3 - 12 • Reactive (not as reactive as Groups 1 or 2), can be free elements • Electron Configuration • ns2(n-1)dxwhere x is column in d-block • Form variable valence state ions • Cations • Examples: Co, Fe, Pt, etc

  16. 2. Nonmetals • Not good conductors • Found on right side of periodic table – AND hydrogen • Usually brittle solids or gases

  17. a. Halogens • Group 17 (7A) • Very reactive • Electron configuration • ns2np5 • Form 1- ions – 1 electron short of noble gas configuration • Anions • Examples: F, Cl, Br, etc

  18. b. Noble Gases • Group 18 • Unreactive, inert, “noble”, stable • Electron configuration • ns2np6full energy level • Have a 0 charge, no ions • Examples: He, Ne, Ar, Kr, etc

  19. 3. Metalloids • Sometimes called semiconductors • Form the “stairstep” between metals and nonmetals • Have properties of both metals and nonmetals • Examples: B, Si, Sb, Te, As, Ge, Po, At

  20. B. Chemical Reactivity • Alkali Metals • Alkaline Earth Metals • Transition Metals • Halogens • Noble Gases

  21. 1A 8A 2A 3A 4A 5A 6A 7A C. Valence Electrons • Valence Electrons • e- in the outermost energy level • Group #A = # of valence e- (except He)

  22. 1A 8A 2A 3A 4A 5A 6A 7A C. Valence Electrons • Valence electrons = • electrons in outermost energy level • You can use the Periodic Table to determine the number of valence electrons • Each group has the same number of valence electrons

  23. D. Lewis Diagrams • Also called electron dot diagrams • Dots represent the valence e- • Ex: Sodium • Ex: Chlorine Lewis Diagram for Oxygen

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