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Chapter 6

Chapter 6. The Periodic Table and Periodic Law. 6.1 Development of the Modern Periodic Table. 1. Antoine Lavoisier (1743-1794) Compiled a list of all known elements 23 know elements at that time. 6.1 Development of the Modern Periodic Table. 2. John Newlands (1837 – 1898)

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Chapter 6

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  1. Chapter 6 The Periodic Table and Periodic Law

  2. 6.1 Development of the Modern Periodic Table • 1. Antoine Lavoisier (1743-1794) • Compiled a list of all known elements • 23 know elements at that time

  3. 6.1 Development of the Modern Periodic Table • 2. John Newlands (1837 – 1898) • Law of Octaves = every 8th element repeats a common set of properties • Not widely accepted due to missing elements and his use of musical terminology

  4. 6.1 Development of the Modern Periodic Table • 3. Dimitri Mendeleev (1834 – 1907) • 1869 he published 1st periodic table by atomic mass and chemical properties • Predicted the properties of missing elements: scandium, gallium, and germanium

  5. 6.1 Development of the Modern Periodic Table • 4. Henry Moseley (1887 – 1915) • 1913 equated number of protons with atomic number • Reordered the P.T. by atomic # fising some of the elements that didn’t fit their spots based on properties

  6. 6.1 Development of the Modern Periodic Table • 5. periodic law: • The repeating pattern of chemical and physical properties when elements are arranged by atomic number

  7. 6.1 Development of the Modern Periodic Table • 6. organization • Groups/families: • Columns of/on the P.T.

  8. 6.1 Development of the Modern Periodic Table • Periods: • Rows of/on the P.T.

  9. 6.1 Development of the Modern Periodic Table • Main group/representative elements: • Elements in groups where the number is followed with an “A” • Have a wide range of chemical and physical properties

  10. TABLE: • Metals: • Location = to the left of the staircase except H • Properties = shiny, mostly solids, good conductors of heat/electricity, and malleable/ductile • Examples = copper (Cu), gold (Au), iron (Fe)

  11. TABLE: • Nonmetals: • Location = to the right of the staircase plus H • Properties = brittle (when solid), mostly gases, poor conductors • Examples =helium (He), oxygen (O), Iodine (I)

  12. TABLE: • Semimetals/metalloids: • Location = along the staircase except Al • Properties = properties of both metals and nonmetals • Examples = B, Si, Ge, As, Sb, Te, Po, and At

  13. TABLE: • Alkali metals: • Location = 1A (except H) • Valence e- and charge = 1 ve- and +1 • Properties = highly reactive; soft, gray solids

  14. TABLE: • Alkaline Earth metals: • Location = 2A • Valence e- and charge = 2 ve- and +2 • Properties = very reactive; soft, gray solids

  15. TABLE: • Transition metals: • Location = “B” groups • Valence e- and charge = 2 ve- and +1,+2,+3 or +4 • Properties = shiny, good conductors, can be polyvalent (means can have more than 1 possible charge)

  16. TABLE: • Halogens: • Location = 7A • Valence e- and charge = 7 ve- and -1 • Properties = highly reactive; can be solids, liquids or gases

  17. TABLE: • Noble Gases: • Location = 8A • Valence e- and charge = 8 ve- and no ion formation • Properties = extremely unreactive, gases, full outer energy level

  18. TABLE: • Rare Earth metals: • Location = Bottom double rows • Valence e- and charge = 2 ve- • Properties = often used as phosphors (elements that emit light when struck by electrons) • Also know as the “Lanthanides” and “Actinides”

  19. 6.2 Classification of the Elements • 1. Valence Electrons = electrons in the highest principle energy level • Within a period: elements have the same # of energy levels as the period # where they are found • Within a group: elements have the same # of ve-s as their group (representative elements) and all transition and rare earth elements have 2 ve-s

  20. 6.2 Classification of the Elements • 2. s, p, d, f blocks • s block: group 1A, 2A, hydrogen and helium • p block: groups 3A-8A but not He • d block: transition elements; all have 2 ve-s because d’s are 1 energy level behind • f block: rare earth elements; all have 2 ve-s because f’s are 2 energy levels behind

  21. 6.3 Periodic Trends • Periodicity: • The repeating nature of the properties of the elements creating common groups (periodic law)

  22. 6.3 Periodic Trends • Atomic Radius • DEFINITION: relative size; distance from the center of the atom to the edge of the e- shell • PERIOD TREND: • GROUP TREND:

  23. 6.3 Periodic Trends • Ionic Radius • DEFINITION: relative size; distance from the center of the ion to the edge of the e- shell • PERIOD TREND: • GROUP TREND:

  24. 6.3 Periodic Trends • Ionization Energy • DEFINITION: the energy required to remove an electron from a gaseous atom • PERIOD TREND: • GROUP TREND:

  25. 6.3 Periodic Trends • Electronegativity • DEFINITION: the ability of an atom to attract electrons while in a chemical bond • PERIOD TREND: • GROUP TREND:

  26. 6.3 Periodic Trends • Lower Left Large (atomic/ionic radius) • Lower Left Low (ionization E./electroneg.)

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