1. The Periodic Table History Structure Trends

2. Part I: Attempts at Classification

3. Dobereiner - 1817 Discovered groups of 3 elements with similar properties “Triads” Ca - 40 Sr – 87.6 Ba - 137 Mass of Sr is about ½ way between Ca and Ba

4. Newlands - 1863 Arranged elements in order of increasing atomic mass Found a repetition of similar properties with every 8 th element “Law of octaves” 7 column table with 7 rows

5. Mendeleev - 1870 Arranged elements in order of atomic mass with 8 long columns and several short rows Arrangement reflected properties of the elements Predicted existence of several elements to fill gaps in his table These elements were later discovered and had the properties predicted

6. Mendeleev’s Periodic Law Properties of the elements are a periodic function of their atomic masses.

7. Moseley - 1913 Performed experiments to determine an accurate mass for several elements which seemed out of place on the table Noticed a pattern in the number of protons Reorganized elements in order of atomic number rather than mass

8. Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers. Why? Patterns of properties repeated

9. Part II: The Modern Periodic Table

10. Periods 7 rows 1 2 3 4 5 6 7

11. Groups or Families 1 2 3456789101112 1314151617 18

12. This Arrangement Reflects: Properties Increasing atomic number Electron configuration

13. Electron Configuration: Li – 1s 2 2s 1 Na – 1s 2 2s 2 2p 6 3s 1 K – 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 All have 1 e- in the outer level All are similar in color and hardness All react vigorously with water Li Na K Same Family

14. C - 1s 2 2s 2 2p 2 N - 1s 2 2s 2 2p 3 O - 1s 2 2s 2 2p 4 All have a full 1 st energy level with e- in s and p orbitals of the second level Number of e- is increasing by 1 CNO Same period

15. Blocks sp d f

16. Classification of Elements: Most of the elements are metals. Located on the left side of the table.

17. General Properties of Metals: 1 to 3 e- in outer energy level of most Lose e- to form bonds (+ ions) Shiny (luster) Hard Good conductors of heat Good conductors of electricity Malleable Ductile Most are solids @ room temp. (Hg)

18. Classification of Elements: Some are nonmetals and are found on the right side of the table.

19. General Properties of Nonmetals: 5 to 8 electrons in the outer level Gain e- to form bond (- ions) Brittle solids or gases Poor conductors of heat Poor conductors of electricity Will share e- w/ other nonmetals to have an octet

20. Classification of Elements: Some elements have properties similar to both metals and nonmetals. These are found bordering the stair- step dividing line. Exception: Al is a metal These elements are called metalloids.

21. Metals, Metalloids, and Nonmetals

22. Some of the Families Have Special Names. 1 – alkali metals 2 – alkaline earth metals 3-12 – transition metals Elements 58 – 71 – lanthanides Elements 90 – 103 – actinides Lanthanides and actinides – rare earth metals 16 – chalcogens 17 – halogens 18 – noble gases

23. Families with special names:

24. Part III: Trends on the Periodic Table

25. Atomic Radius Radius of an atom without regard to surrounding atoms (size) Radius depends on: the number of energy levels the strength of the nucleus r

26. As n increases, size of the e- cloud increases. Each period represents a higher principal quantum number (n) F Cl Br Atomic radius increases down a family.

27. Across a period, nuclear charge increases by 1 for each element. A stronger nucleus acts like a stronger magnet which attracts the e- cloud. CON Atomic radius decreases across a period.

28. General Trend for Atomic Radius: Increases

29. Confirming the Trend: Look at the following graph of actual data. Notice the pattern for elements in the same period. Notice the pattern for elements in the same family.

30. Atomic Radius

31. Use position in the periodic table to determine which is larger? Na or Rb? Cl or I? Al or Si? K or Ca? Ag or Au? Ni or Cu? La or U? H or He?

32. Ionic Radius Size of an ion Ions are charged atoms formed when: Atoms lose e-  (+ Ion)  Cation Atoms gain e-  (- Ion)  Anion

33. Cations are smaller than their respective neutral atoms. Metals usually lose all valence electrons. The nucleus pulls tighter on the remaining electrons. + - e- + + more p+ than e-

34. Anions are larger than their respective neutral atoms. Nonmetals usually gain electrons to complete the valence shell. Electrons repel each other and spread out more. + e- more e- than p+ - - -

35. Which is larger? Ca or Ca +2 F or F –1 K or K +1 O or O -2 F -, Ne or Na +

36. First Ionization Energy Energy needed to remove the most loosely held electron from a gaseous atom. Factors that affect ionization energy: Radius Nuclear charge Shielding effect Stability of sublevels

37. Radius: The greater the distance between the nucleus and the valence electrons, the easier it is to lose an electron.

38. Nuclear charge: Within a period, the higher the nuclear charge, the higher the ionization energy.

39. Shielding Effect: Other e- block the pull of the nucleus on the outer e-. Electrons repel each other.

40. Stability of Sublevels: e- in filled or half-filled sublevels are extremely hard to remove. Higher ionization energy than their immediate neighbors CNO Highest IE

41. General Trend for IE: Ionization energy increases as you go up a family and across a period.

42. Confirming the Trend: Look at the following graph of actual data. Notice the pattern for elements in the same period. Notice the pattern for elements in the same family.

43. First Ionization Energies

44. General Trend for Ionization Energy: Increases

45. Electron Affinity: Attraction of an atom for an additional electron. Factors affecting EA: Size Nuclear charge Shielding effect Stability of sublevels

46. Size: In large atoms, the nucleus exerts less pull on the outer level.

47. Nuclear charge: Within a period, an increase in nuclear charge creates a greater attractive force.

48. Shielding effect: High numbers of electrons repel additional electrons

49. Stability of sublevels: Filled or half-filled sublevels are more stable. Don’t need any more electrons

50. General Trend for EA: EA increases as you go up a family and across a period.

51. Confirming the Trend: Look at the following graph of actual data. Notice the pattern for elements in the same period. Notice the pattern for elements in the same family.

52. Electron Affinity:

53. First 2 periods: Filled sublevel

54. First 2 periods: Half-filled sublevel

55. General Trend for Electron Affinity: Increases

56. Oxidation Numbers: Position in the periodic table can be used to predict oxidation numbers.

57. Elements with e- configurations ending with s x : Lose X e- to form cations

58. Elements with e- configurations ending with d x : Lose s e- first Lose d e- one at a time  If more than 5 d e-, will lose only those in excess of 5  If 10 d e-, will lose only s e-

59. Elements with e- configurations ending with p x : Gain e- to complete p and form anions Lose s and p e- to form cations Lose only p e-

60. Examples: Ca – 4 s 2 + 2 V – 4s 2 3d 3 + 2, + 3, + 4 and + 5 Fe – 4s 2 3d 6 +2 and +3 O – 2s 2 2p 4 - 2, +6 and +4

61. Other Trends: Look at the following graphs to determine if there is a trend for each of these properties.

62. Melting Points of the Elements :

63. Density:

64. Ionization Energy: Group Trend As the number of energy levels increase, shielding was greater. Shielding is interference from having energy levels in b/w the nucleus and the valence e. It takes less energy to get an e from an atom with multiple energy levels.

65. General Trend for Electronegativity: Increases

66. Electronegativity: Ability to draw an e from an element Best e stealer: F (=4) Two elements closest to F, O and Cl, are the next best Group: The fewer energy levels there are, the more the nucleus can pull and attract an e - Period: As you add more protons in the nucleus, there is a bigger attraction for e -.

67. Electronegativity: