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Exploring the Periodic Table: History, Laws, and Patterns

Learn about the evolution of the periodic table, from Joseph Proust to Modern Periodic Law, and understand the arrangement, laws, and properties of elements. Discover the significant contributions of scientists and the patterns found in the table.

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Exploring the Periodic Table: History, Laws, and Patterns

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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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