1. Chapter 6 The Periodic Law http://www.privatehand.com/flash/elements.html

3. History of the Periodic Table 1869 – Dmitri Mendeleev published his periodic table. He had arranged it by grouping together the elements that had similar properties and by increasing atomic masses. His periodic table left empty spaces for new elements that would be discovered.

4. Mendeleev’s List of elements in Russian Circa 1869

5. Periodic Table in English (Circa 1891)

6. Periodic Table circa 1898

7. History of the Periodic Table u 1911 – Henry Moseley (a student of Ernest Rutherford) rearranged a few elements on the periodic table so that elements were arranged by increasing atomic number rather than by atomic mass.

8. History of the Periodic Table u 1944 – Glenn T. Seaborg rearranged the periodic table to make it look like it does today. He moved the Actinide Series and the Lanthanide Series elements to the bottom of the periodic table. http://livingtextbook.or egonstate.edu/media/vi d/lbl5a3.mov

9. Periodic Table Circa 1944

10. Modern Periodic Table

11. Parts of the Periodic Table u The periodic table can be divided and labeled using several methods.

12. Elements are arranged: Vertically columns are called Groups Horizontal rows are called Periods Parts of the Periodic Table

13. Periodic Families u Alkali Metals u Alkaline Earth Metals u Halogens u Noble gases

15. Parts of the Periodic Table u Metals u Non-metals u Metalloids

17. Parts of the Periodic Table u Main Group or Representative Elements u Transition Metals u Rare earth elements u Trans Uranium Elements

19. The Periodic Law u The physical and chemical properties of the elements are periodic functions of their atomic numbers

21. Periodic Trends If you understand the trends on the periodic table, you can predict almost anything about any element on the periodic table. We will study: Atomic Radii Ionic Radii Valence Electrons Reactivity Electronegativity Electron Configuration

22. Atomic Radius u Define – One-half the distance between the nuclei of identical atoms that are bonded together. Periodic Trend: The atomic radius decreases as you go across the periodic table and increases as you go down the periodic table

23. Periodic Trends u As you go across a period, the radius gets smaller. u Electrons are in same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

24. Group trends u As we go down a group... u each atom has another energy level, u so the atoms get bigger. H Li Na K Rb

25. Trends in Ionic Size u Cations form by losing electrons. u Cations are smaller than the atom they come from. u Metals form cations. u Cations of representative elements have noble gas configuration.

26. Ionic size u Anions form by gaining electrons. u Anions are bigger than the atom they come from. u Nonmetals form anions. u Anions of ‘main’ groups elements have noble gas configuration.

27. Group trends u Adding energy level u Ions get bigger as you go down. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

28. Periodic Trends u Across the period, nuclear charge increases so they get smaller. u Energy level changes between anions and cations. Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1-

29. Valence Electrons u Define: The electrons available to be lost, gained, or shared in the formation of compounds. u The electrons in the highest energy level

30. Valence Electrons Periodic Trends: Group 1 = 1 valence electron = 1+ Oxidation Number Group 2 = 2 valence electrons = 2+ Oxidation Number Group 13 = 3 valence electrons = 3+ Oxidation Number Group 14 = 4 valence electrons = 4+/4- Oxidation Number Group 15 = 5 valence electrons = 3- Oxidation Number Group 16 = 6 valence electrons = 2- Oxidation Number Group 17 = 7 valence electrons = 1- Oxidation Number Group 18 = 8 valence electrons = 0 Oxidation Number

31. 1+ H Li Na Be 2+ Mg B 3+ Al C 4+ 4- Si N 3- O 2- F 1- Ne 0 He PSCl K Rb Cs Fr Ca Sr Ba Ra Ga In Tl Ge Sn Pb AsSe Ar BrKr SbTeIXe BiPoAtRn

32. Reactivity u Reactivity increases as you go down the columns of metallic elements. u Reactivity decreases as you go down the columns of non-metallic elements. u Watch the video to see what that means.

33. Electronegativity u The tendency for an atom to attract electrons to itself when it is chemically combined with another element. u High electronegativity means it pulls the electron toward it.

34. Group Trend u The further down a group, the farther the electron is away, and the more electrons an atom has. u More willing to share. u Low electronegativity.

35. Periodic Trend u Metals are at the left of the table. u They let their electrons go easily u Low electronegativity u At the right end are the nonmetals. u They want more electrons. u Try to take them away from others u High electronegativity.

36. Electronegativity (important to determine bond type)

37. Electron Configuration and the Periodic Table

38. Electron Configuration of Main Group Elements Group Period #(+) Example 1s 1 Na = 3s 1 2s 2 Ba = 6s 2 13s 2 p 1 Ga = 4s 2 4p 1 15s 2 p 3 Sb = 5s 2 5p 3 17 18 s2p5s2p5 Br = 4s 2 4p 5 s2p6s2p6 Rn =6s 2 6p 6

39. Electron Configuration Transition Elements Period # s 2 + Period # (-1) d 1 – 10 Examples: Sc = 4s 2 3d 1 Zn = 4s 2 3d 10 Mo = Ir = 5s24d4 6s25d7

40. Electron Configuration for Lanthanide and Actinide Series Period # s 2 + Period # (- 2) f 1 - 14 Examples: Ce = 6s 2 4f 1 Ho = 6s 2 4f 10 U = Bk = 7s25f3 7s25f8

41. Atomic size increases, shielding constant Ionic size increases

42. Ionization energy, Electronegativity, and Electron Affinity INCREASE