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The Periodic Table

The Periodic Table. Objectives: describe the origin of the periodic table state the periodic law explain the relationship between electron configurations and the location of elements in the periodic table describe the nature of periods and groups of elements in the periodic table

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The Periodic Table

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  1. The Periodic Table Objectives: describe the origin of the periodic table state the periodic law explain the relationship between electron configurations and the location of elements in the periodic table describe the nature of periods and groups of elements in the periodic table state the definitions of some properties of the elements that exhibit periodicity describe the trends of those properties within periods and groups of elements Created by C. Ippolito Nov. 2006

  2. Early Classification Systems • 1817 Johann Dobreiner • noticed that Ca, Ba, and Sr possess similar properties • noticed that mass of Sr was halfway between Ca and Ba • called this “group” a triad • found two additional triads • Cl, I and Br • S, Te, and Se • 1863 John Newlands • arranged elements in atomic mass order • properties repeat every eighth element • Law of Octaves • made seven groups of seven elements each Created by C. Ippolito Nov. 2006

  3. Mendeleev’s Periodic Table • 1869 Dmitri Mendeleev • properties were a function of atomic mass • periods of varying lengths 7, 7, 17 • left blank spots so he could put similar properties in same column • predicted properties of “empty” spots • when discovered they matched his predictions • Mendeleev’s Periodic Law • the properties of elements repeat in an orderly fashion as a function of their atomic masses Created by C. Ippolito Nov. 2006

  4. Modern Periodic Law • Mendeleev problems • I and Te – I before Te by mass; I after Te by properties • put them by properties; thought masses measured incorrectly • Co and Nickel; K and Ar similar problems • Henry Moseley • properties of elements are a periodic function of their atomic number not atomic mass Created by C. Ippolito Nov. 2006

  5. Reading the Periodic Table • Periods • horizontal rows • begin on left with a metal and end on right with noble gas • Family or Group • vertical rows • similar physical and chemical properties • 1 – 2 and 13 – 18 • have same number of valence electrons • 3 – 12 • s level of outermost energy level has 1 or 2 electrons • d levels are “being” filled Created by C. Ippolito Nov. 2006

  6. Periods of Elements • Period Number • the energy level of its Valence Electrons • Short Periods (1, 2 and 3) • elements with up to 2e-, 8e- and 8e- • Period 1 (H and He) • completes 1s level • Period 2 (Li, Be, B, C, N, O, F, and Ne) • begins filling 2nd Level with Ne having 2s and 2p levels completed • Period 3 (Na, Mg, Al, Si, P, S, Cl and Ar) • begins filling 3rd Level with Ar having 3s and 3p levels completed Created by C. Ippolito Nov. 2006

  7. Periods of Elements • Period Number • the energy level of its Valence Electrons • Long Periods (4 and 5) • each contain 18 elements • Period 4 (K thru Kr) • K and Ca are completing 4s level • Sc thru Cu completed 4s and are filling 3d level • these elements are the transitional metals or elements • Ga thru Kr are completing the 4p level • Period 5 (Rb thru Xe) • Rb and Sr are completing 5s level • Y thru Cd completed 5s and are filling 4d level • these elements are the transitional metals orelements • In thru Xe are completing the 5p level Created by C. Ippolito Nov. 2006

  8. Periods of Elements • Period Number • the energy level of its Valence Electrons • Long Periods (6 and 7) • each contain 32 elements • Period 6 (Cs thru Rn) • Lanthanides • adding electrons to 4f level • Period 7 (Fr thru Uuo) • Actinides • adding electrons to the 5f level Created by C. Ippolito Nov. 2006

  9. Groups of Elements • Group 1 – Alkali Metals • lose 1e- to form +1 ions • Group 2 – Alkaline Earth Metals • lose 2e- to form +2 ions • Group 13 – no special name • lose 3e- to form +3 ions • Group 14 – no special name • lose 4e- to form +4 ions • gain 4e- to form -4 ions • Group 15 – Nitrogen Family • gain 3e- to form -3 ions • Group 16 – Oxygen Family • gain 2e- to form -2 ions • Group 17 – Halogens • gain 1e- to form -1 ions • Group 18 – Noble Gases • complete valence level Created by C. Ippolito Nov. 2006

  10. Periodic Table of Elements • metals • metalloids • nonmetals Created by C. Ippolito Nov. 2006

  11. Review of the Periodic Table • Metals are on left side; Nonmetals on right side • Majority of elements are metals • Each period “opens” a new energy level with a different principal quantum number • Each period represents atoms with “larger” electron clouds Created by C. Ippolito Nov. 2006

  12. Periodicity in Properties • Properties directly related to the attraction of positive nucleus for negative electrons • Coulombic attraction • depends on both the quantity of charge and distance separating charges • Ionization Energy • Electronegativity • Number of Electrons • Atomic Radius • Ionic Radius Created by C. Ippolito Nov. 2006

  13. Periodicity of Ionization Energy • Ionization Energy • removes the most loosely held electron from the outer energy level of an atom in the gas phase. • Periodicity • ionization energy increases as atomic number increases (left to right) in a period • metals have LOW IONIZATION ENERGIES • nonmetals have HIGH IONIZATION ENERGIES • ionization energy decreases as atomic number increases (top to bottom) in a group • Shielding Effect • inner electrons block attraction of nucleus to outer electrons • kernel electrons repel valence electrons Created by C. Ippolito Nov. 2006

  14. Ionization Energy vs Atomic Number Created by C. Ippolito Nov. 2006

  15. Factors affecting Ionization Energy • Nuclear Charge • increase in charge increases ionization energy • Shielding • increase in shielding decreases ionization energy • Radius • increase in radius decreases ionization energy • Sublevel • full or half full levels require additional energy to be removed Created by C. Ippolito Nov. 2006

  16. Electron Affinity vs. Electronegativity • Electron Affinity • ability of an atom’s nucleus to attract additional electrons • as ionization energy  electron affinity  • metals have LOW electronegativity • nonmetal have HIGH electronegativity • Electronegativity • comparative scale relating element’s atoms ability to attract electrons when bonded • increases from left to right in a period; decreases from top to bottom of a family • high bonds with low, the greater the difference the stronger the bond Created by C. Ippolito Nov. 2006

  17. Atomic Radius and Periodicity • Atomic Radius • closest distance one atom can approach another atom • Covalent Radius • distance from nucleus to valence shell when covalently bonded • decreases left to right in a period; increases top to bottom in a family • van der Waals Radius • distance between nuclei of identical unbonded atoms • Ionic Radius • distance from nucleus to valence electrons in a monoatomic ion • no clear pattern in period; increases top to bottom in family Created by C. Ippolito Nov. 2006

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