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Periodic Table

Periodic Table. Chapter 6. What do I know?. On the back of the blank periodic table write down at least 3 pieces of information you can get from the periodic table. A Brief History…. Joseph Proust Law of Definite Composition elements combine in definite proportions by weight

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Periodic Table

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  1. Periodic Table Chapter 6

  2. What do I know? • On the back of the blank periodic table write down at least 3 pieces of information you can get from the periodic table.

  3. A Brief History… • Joseph Proust • Law of Definite Composition • elements combine in definite proportions by weight • The weight of one element that combines with the weight of another element = combining weight

  4. Joseph Berzelius [1807 - 1818 ] Determined the combining weights of 43 elements with oxygen. Recognized similarities of certain elements... similar metallic properties similar reactive properties Li, Na, K similar nonmetals Cl, Br, I, “TRIPLETS”

  5. Johann Wolfgang Dobereiner • 1829 • mathematician • discovered that combining weight of middle triplet is the average [or near average]of the combining weights of the other two • Li, Na, K

  6. Jean Stas • 1860 • confirmed Proust theory of definite composition • established accurate atomic weight of the known elements

  7. Was there a relationship between the weight of an element and its properties? • John A.R. Newlands • 1865 • arranged elements in order of atomic weight • elements with similar properties were 7, or multiple of 7 apart • Law of Octaves

  8. Dimitri Mendeleev 1869 • developed a chart -listed elements by increasing atomic weight • grouped elements with similar properties in the same row • Left Gaps where no element fit the pattern. • Predicted discovery of new elements • Predicted properties of new elements

  9. Mendeleev’s Table

  10. . “The properties of elements are in periodic dependence of their atomic weights.” Dimitri Mendeleev Old Periodic Law

  11. ALTERATIONS and ADDITIONS • Sir Wm.Ramsay • 1890’s • Discovered Ne, Ar, Kr, Xe • Helium and Radon disc. Previously • New row added to Periodic Table

  12. Henry Gwyn-Jeffreys Mosley • 1914-1915 • Number of protons determined • atomic number - identifies what an element is • Periodic Table Rearranged • elements arranged by increasing atomic number • similar elements put in columns instead of rows

  13. Modern Periodic Law • “The properties of elements are in periodic dependence of their atomic numbers.”

  14. ARRANGEMENT OF THE MODERN PERIODIC TABLE • A horizonal row on the periodic chart is refered to as either a period, or a series. • A vertical column on the periodic chart is refered to as either a group, or a family.

  15. Element Location Electron Dot Properties on Chart Notation • H Grp 1 H .Colorless gas • LiGrp 1 Li . Soft; silver highly reactive • Na Grp 1 Na . Soft; silver highly reactive • K Grp 1 K . Soft; silver highly reactive • Rb Grp 1 Rb . Soft; silver highly reactive • Cs Grp 1 Cs . Soft; silver highly reactive • Fr Grp 1 Fr . Soft; silver • most reactive metal Alkali Metals

  16. Element Location Electron Dot Properties on Chart Notationc • Be Grp 2 Be : Reactive metal • Mg Grp 2 Mg : Reactive metal • Ca Grp 2 Ca : Reactive metal • Sr Grp 2 Sr : Reactive metal • Ba Grp 2 Ba : Reactive metal • Ra Grp 2 Ra : Most reactive metal of group Alkaline Earth Metals

  17. What pattern(s) do we see? • All elements in groups have same electron dot structure. • Group placement predicts valence. • Groups usually have similar properties. • Most reactive metals at the bottom of the group.

  18. Element Location Electron Dot Properties on Chart Notation • B Grp 3 B : nonmetal; black solid • Al Grp 3 Al: Metal • Ga Grp 3 Ga : Metal • In Grp 3 In : Metal • Tl Grp 3 Tl : Most reactive metal

  19. Element Location Electron Dot Properties on Chart Notation • C Grp 4 C : black→clear solid • Si Grp 4 Si : Metalloid • Ge Grp 4 Ge : Metal • Sn Grp 4 Sn : Metal • Pb Grp 4 Pb : Most reactive metal

  20. Element Location Electron Dot Properties on Chart Notation • N • Grp 5 N : gas; nonmetal • P • Grp 5 P : nonmetal • As • Grp 5 As : Metalloid • Sb • Grp 5 Sb : Metalloid • Bi • Grp 5 Bi : Metal

  21. Element Location Electron Dot Properties on Chart Notation • O • Grp 6 O : gas; nonmetal reactive • S • Grp 6 S : Nonmetal • Se • Grp 6 Se : Nonmetal • Te • Grp 6 Te : Nonmetal • Po • Grp 6 Po : Metal Chalcogen Family

  22. Element Location Electron Dot Properties on Chart Notation • F • Grp 7 : F : gas; most reactive nonmetal • Cl • Grp 7 :Cl : gas; reactive nonmetal • Br • Grp 7 :Br : liquid; reactive nonmetal • I • Grp 7 : I : solid; reactive nonmetal • At • Grp 7 :At : solid; reactive nonmetal Halogen Family

  23. Element Location Electron Dot Properties on Chart Notation • He • Grp 8 He : inert; nonmetal • Ne • Grp8 :Ne : inert; nonmetal • Ar • Grp 8 :Ar : inert; nonmetal • Kr • Grp 8 :Kr : inert; nonmetal • Xe • Grp 8 :Xe : inert; nonmetal • Rn Grp 8 :Rn : inert; nonmetal Noble Gases / Inerts

  24. What pattern(s) do we see? • All elements in groups have same electron dot structure. • Group placement predicts valence. • Groups usually have similar properties – (exception: steps) • Most reactive nonmetals at the top of the group. • Most reactive metals at the bottom of the group.

  25. I spy with my little eye an element with… • 3 energy levels and 2 valence electrons • Mg • 5 energy levels and 4 valence electrons • Sn • 2 energy levels and 8 valence electrons • Ne • 1 valence electron and 5 energy levels • Rb • 1 valence electron and 7 energy levels • Fr

  26. I spy with my little eye an element with… • 4 energy levels and 7 valence electrons • Br • 3 energy levels and 5 valence electrons • P • 2 valence electrons and 4 energy levels • Ca • 3 valence electrons and 2 energy levels • B • 8 valence electrons and 5 energy levels • Xe

  27. I spy with my little eye an element with… • The heaviest halogen… • At (astatine) • The triplet with the average atomic weight of 35.5… • Cl • The least reactive Chalcogen • Po (polonium) • The group that fills the s2 valence orbital • Alkaline Earth Metals • A third period metalloid • Si

  28. Bonding • See interactive

  29. Types of Bonding • Ionic • Electrons transfer from one atom to another creating + and – ions. • Covalent • Atoms share electrons to create a molecule. • Metallic • Many atoms share electrons

  30. Types of Bonding • Ionic • Electrons transfer from one atom to another • creating + and – ions. e- + - + energy

  31. Ionization Energy • The energy required to remove the outermost e- in an atom. Helium Neon Argon Hydrogen Lithium Sodium

  32. Why are some e- removed more easily? • Electrons that are farther away from the nucleus and that have more E levels between them and the nucleus • Low ionization energy • characteristic of METALS. • High ionization energy • characteristic of NONMETALS. • Removing successive electrons is more difficult, but follows the same overall pattern. • Na + Energy  Na+ + e- 119 Kcal / mol • Na+ + Energy  Na++ + e- 1090 Kcal/ mol • Na+++ Energy  Na++++ e- 1652 Kcal/ mol

  33. Electron Affinity The energy released / absorbed when an electron is accepted by a neutral atom e- + - + energy Ionization E removes e- and forms + ion linked Electron affinity is the E released when the neutral atom accepts the freed e- and becomes -

  34. Electron Affinity Lithium Sodium Fluorine Chlorine

  35. Electron Affinity Increases across a period Decreases

  36. For atoms that have - valences: • Atom + e-  A- + E • exothermic - energy released • (electron affinity) • stable product • Atom + e- + E  A- • endothermic - energy required • unstable product

  37. Covalent • Atoms share electrons to create a molecule. Shared e-’s

  38. Electronegativity • the attraction of an atom for a shared pair of electrons

  39. Electronegativity Fluorine Chlorine Lithium Sodium

  40. Table of Electronegativities

  41. Electronegativity • Types of Covalent Bonds: • pure covalent - relatively even sharing of e- • polar covalent - uneven sharing of e- • 0 - .5 ....... pure covalent • .5 - 1.7..... polar covalent • > 1.7 ....... ionic bond

  42. Atomic Radius [size] Sodium Lithium Chlorine Fluorine

  43. Atomic Radius [size]

  44. Down a group • E levels are added. • Across a period • Increased attraction between the E levels and the nucleus causes the size to decrease. • Pauli Repulsion Theory • As the number of electrons increases so does the repulsion between the electrons; this may help account for the irregular increase in the radii.

  45. Ions [size] • Increases down a Group • Decreases across a Period • Metal atoms lose electrons • become positive (cation) • Cations are SMALLER than the atoms from which they come. • Nonmetal atoms gain electrons • become negative (anion) • Anions are LARGER than the atoms from which they come.

  46. Density Aluminum Boron Sodium Chlorine Lithium Fluorine

  47. Density • Here the density of each period is graphed individually

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