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The Periodic Table

The Periodic Table. By 1860 60 elements had been discovered but scientists had no way of organizing them. J.W. Dobereiner classified elements that had similar properties into triads, organizing them by atomic mass.

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The Periodic Table

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  1. The Periodic Table

  2. By 1860 60 elements had been discovered but scientists had no way of organizing them

  3. J.W. Dobereiner classified elements that had similar properties into triads, organizing them by atomic mass

  4. Triads were useful because they grouped elements with similar properties and revealed an orderly pattern in some of their physical and chemical properties which are related to related to atomic mass.

  5. Dimitri Mendeleev • 1869 – grouped elements by atomic mass • Chemical and physical properties repeated in an orderly way when elements are arranged according to increasing atomic mass • Elements with similar properties were placed in horizontal rows • Density, melting point, and boiling point increase as atomic mass increases

  6. Elements on the periodic table show periodicity • The tendency to recur at regular intervals • Mendeleev was able to predict the existence and properties of elements that had not yet been discovered

  7. Henry Moseley • Rearranged the elements according to increasing atomic number • Atomic # = • Resulted in the structure of the modern periodic table

  8. Periodic Law – the physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number.

  9. Periods and Groups • 7 horizontal rows = periods • Correspond to outermost energy level • Vertical columns = groups/families • Correspond to the number of outermost electrons • Have similar properties • Some have special names

  10. Group/Family Names • Group 1 = Alkali metals • Group 2 = Alkaline earth metals • Group 3 – 12 = Transition metals • Inner Transition metals • Group 17 (VII A) = Halogens • Group 18 (VIII A) = Nobel gases

  11. Physical States and Classes of Elements • Most elements are solid at room temperature • Br & Hg are liquid • N, O, F, Cl, and Noble gases are gas

  12. Elements are classified as metals, metalloids, or nonmetals based on their properties

  13. Metals • Located to the left of the stairs • Have luster • Conduct heat and electricity • Usually bend without breaking • Solid at room temperature • Very high melting points

  14. Nonmetals – located to the right of the stairs • Brittle • Dull looking • Poor conductors of heat and electricity • Usually gases

  15. Metalloids – have properties of both metals and nonmetals • B, Si, Ge, As, Sb, Te, Po, At • Some are semiconductors – conducts electricity better than a nonmetal but not as good as a metal

  16. Patterns in Valence Electrons • Valence electron = • All elements of a family have the same number of valence electrons • Increase across a period

  17. There is a relationship between the electron configuration of an element and its placement on the periodic table

  18. Alkali Metals (1A) Na K

  19. Alkaline Earth Metals (2A) Mg Ca

  20. Group 6A O S

  21. Halogens (7A) F Cl

  22. Nobel Gasses (8A) Outermost s & p sublevels completely filled (non-reactive) Filled outer levels make atoms stable and non-reactive

  23. Representative Elements Outermost s & p sublevels are only partially filled (group A elements) Group # = # of valence electrons

  24. 1A = s1 (1 valence electron) 2A = s2 (2 valence electrons) 3A = s2p1 (3valence electrons) 4A = s2p2 (4 valence electrons) 5A = s2p3 (5 valence electrons) 6A = s2p4 (6 valence electrons) 7A = s2p5 (7 valence electrons) 8A = s2p6 (8 valence electrons)

  25. Transition Metals Outermost s and nearby d sublevels contain electrons All transition metals usually have 1, 2, or 3 valence electrons

  26. Inner Transition Metals Outermost s and nearby f sublevels contain electrons

  27. Movement of Electrons 1A = s1 (1 ve, easily lost) 2A = s2 (2 ve, easily lost) 3A = s2p1 (3 ve, easily lost) 4A = s2p2 (4 valence electrons) 5A = s2p3 (5 ve, 3 gained) 6A = s2p4 (6 ve, 2 gained) 7A = s2p5 (7 ve, 1 gained) 8A = s2p6 (8 valence electrons)

  28. Why do atoms form ions? Representative elements lose or gain electrons in order to obtain the same electron configuration as a noble gas

  29. Before • Na • 1s22s22p63s1 • B • 1s22s22p1 • P • 1s22s22p63s23p5 • F • 1s22s22p5

  30. After • Na1+ • 1s22s22p6 (Ne) • B3+ • 1s2 (He) • P3- • 1s22s22p63s23p6 (Ar) • F1- • 1s22s22p6 (Ne)

  31. Charges of Ions 1A = 1+ 2A = 2+ 3A = 3+ 5A = 3- 6A = 2- 7A = 1-

  32. Trends in Ionization Energy and Electron Affinity Ionization energy = amount of energy needed to remove an electron from an atom (J) Increases as you go UP a group Increases as you go across a period (left to right)

  33. Ionization Energy cont. • Increases up a group • Hard to remove electrons close to the nucleus • Increases across a period • Hard to remove electrons from small atoms

  34. Electronegativity Tendency for the atoms of the element to attract electrons when they are chemically combined with another element (form a negative ion) Each atom is assigned an electronegativity value

  35. Increases up a group • Closer outer energy levels attract electrons more • Increases across a period (left to right)

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