1. Periodic Relationships Among the Elements Chapter 5 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. History of the Periodic Table

3. Arranged the elements by increasing ATOMIC MASS and saw a periodic repetition of properties Produced the first PERIODIC TABLE – 1871 The table placed elements with similar properties in the same column Kept “holes” for undiscovered elements, and predicted the properties in advance Dimitri Mendeleev

4. Mendeleev’s Periodic Table

5. Properties of elements predicted by Mendeleev

6. H.G. Moseley in 1914 Rearranged the elements by: ATOMIC NUMBER This has become the MODERN PERIODIC TABLE

7. Electrons and Ions on the Periodic Table

8. Review: Valence Electrons What are valence electrons? *Remember: Elements in a group have similar properties because they have the same valence electron configuration

9. Valence Electron Configuration Groupe- configValence electrons Expected Charges 1ns 1 2ns 2 13ns 2 np 1 14ns 2 np 2 15ns 2 np 3 16ns 2 np 4 17ns 2 np 5 18ns 2 np 6

10. +1+2+3 -2-3 Charges Of Representative Elements 8.2

11. Na + Al 3+ F-F- O 2- N 3- What ions are isoelectronic with Neon? Isoelectronic: Elements and ions that have the same number of electrons and therefore the same electron configuration Mg 2+ What would the electron configuration be?___________

12. When a cation is formed from an atom of a transition metal, electrons are always removed first from the s orbital and then from the d orbitals. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Forming Ions with Transition Metals

13. Periodic Table Groups and Properties

14. Periodic Table Groups

15. Very unreactive due to full valence shell Odorless, colorless, gases

16. 1

17. Extremely reactive; not found in pure form in nature Silverly, soft form strong bases in water

19. less reactive than alkali, but still not found in pure form also form strong bases in water

20. Transition Metals typical metallic properties form colorful ions

22. shiny metals similar reactivity to Group 2 filling 4f sublevel

24. mainly synthetic (created in a lab) all radioactive

26. very reactive volatile (exist as gases)

27. 1

28. Properties of Metals 1.shiny (luster) 2.conductors of heat and electricity 3.reactive with acids 4.ductile – can be stretched into a wire 5. malleable –can be hammered or rolled into sheets 6. forms positive ions (by losing e - )

29. Properties of Nonmetals 1.dull and brittle 2.poor conductors of heat and electricity 3.does not react with acids 4.usually gases at room temp. 5.forms negative ions (by gaining e - )

30. What are properties of Metalloids?? In the middle! Metalloids have properties of BOTH!! (metals and nonmetals)

31. Periodic Trends

32. PERIODIC LAW When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic (repeating) pattern. patterns on the periodic table are called periodic trends

33. Atomic Radius half the distance from center-center of 2 like atoms

34. Atomic Radii DOWN a Group ↓ As you go down there are more energy levels, the atom size gets larger ↓ There are more electrons between the nucleus and the outermost energy level which increases the shielding effect

35. Shielding Effect reduction of attraction between positive nucleus and outermost electrons outer electrons are not held tight and can move away

36. Na Atomic Radius: down group P P P X X X P X P X P X P P P P P X X X X X X P

37. K P P P X X X P X P X P X P P P P P X X X X X P X X X X X X X X X

38. Atomic Radii DOWN a Group ↓ DOWN THE GROUP ATOMIC RADIUS INCREASES more energy levels, the larger the size of the atom

39. Atomic Radii ACROSS a Period → Each atom gains one proton and one electron in the same energy level → Each added electron is the same distance from the nucleus → The positive charge increases and exerts a greater force on the electrons pulling them closer to the nucleus

40. REMEMBER! PROTONS are bigger and stronger! electrons are smaller and weaker! P ++ - - e

41. Ask yourself, how effective are the positive protons pulling in the electrons? Atomic Radii ACROSS a Period Effective nuclear charge: “positive charge” felt by an electron. Within a period, every time a proton is added, the effective nuclear charge increases… so the radius decreases

42. Atomic Radius: across period P P P X X X P X P X P X P P P P P X X X X X

43. Atomic Radii ACROSS a Period → ACROSS THE PERIOD ATOMIC RADIUS DECREASES greater effective nuclear charge (more protons), greater pull on the electrons, smaller radius

44. Ionic Radii half the distance from center- center of 2 like ions

45. Ionic Radius DOWN a Group ↓ As you go down a group another energy level is added, increasing the size of the atom. (just like the atomic radius)

46. Ionic Radius DOWN the Group ↓ DOWN THE GROUP IONIC RADIUS INCREASES more energy levels, increase in atom size

47. Ionic Radius ACROSS the Period Cation: positive ion formed from losing an electron → A cation is always smaller than the original atom → The more electrons lost the more protons available to attract a smaller number of electrons.

48. Na + Ionic Radius P P P X P X X P P P P P P P X X X X X X X X X P

49. Ionic Radius ACROSS the Period → ACROSS THE PERIOD IONIC RADIUS DECREASES greater effective nuclear charge, less electrons, the shorter the radius

50. Ionic Radius ACROSS the Period Anion: negative ion formed from gaining an electron → A anion is always larger than the original atom → The more electrons gained, the less protons available to attract a larger number of electrons.

51. F Ionic Radius P P P X P X X P P P P P P P X X X X X X X X P -

52. Ionic Radius ACROSS the Period → ACROSS THE PERIOD IONIC RADIUS DECREASES As electrons are added the atom gets larger from right to left, General trend from left to right is decreasing

53. Ionic Radii

54. amount of energy needed to remove an electron from an atom Ionization Energy

55. X X + + e - X X 2 + + e - X X 3 + + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy I 1 < I 2 < I 3 Multiple Ionization Energies

56. Ionization Energy DOWN a Group ↓ As you go down a group atoms become larger ↓ The more electrons in an atom between the nucleus and valence shell, the greater the shielding effect

57. Ionization Energy DOWN a Group ↓ DOWN THE GROUP IONIZATION ENERGY DECREASES greater distance from the nucleus, greater shielding effect less energy needed to remove electron

58. Ionization Energy ACROSS a Period → As atomic radius decreases there is a greater attraction between protons and electrons. (effective nuclear charge) → The stronger the attraction, the more energy needed to remove an electron.

59. Ionization Energy ACROSS a Period → ACROSS THE PERIOD IONIZATION ENERGY INCREASES greater the effective nuclear charge, more energy required to remove electron

60. Electronegativity It is a “tug of war” between the two atoms of a bond H F : : : Which is the more electronegative element?... ability of an atom to attract electrons

61. Electronegativity DOWN the Group ↓ The farther away from the nucleus, the greater the shielding effect ↓ The larger the atom, the less likely it is to accept more electrons.

62. Electronegativity DOWN the Group ↓ DOWN THE GROUP ELECTRONEGATIVITY DECREASES farther the distance from the nucleus, lower ability to attract electrons

63. Electronegativity ACROSS the Period → As you go across a period atomic radius decreases because there is a greater effective nuclear charge → Metals do not attract electrons. → Non-metals do attract electrons.

64. Electronegativity ACROSS the Period → ACROSS THE PERIOD ELECTRONEGATIVITY INCREASES greater effective nuclear charge, greater ability to attract electrons

65. the energy change that occurs when an electron is added to an atom to form an anion. Increases with ability to attract and hold an electron (electronegativity) Electron Affinity

66. Electron Affinity DOWN the Group ↓ The larger the atom the more difficult to accept electrons

67. Electron Affinity DOWN the Group ↓ DOWN THE GROUP ELECTRON AFFINITY DECREASES farther the distance from the nucleus, does not want to gain electrons

68. Electron Affinity ACROSS the Period → As effective nuclear charge gets stronger, it is easier to attract an electron.

69. Electron Affinity ACROSS the Period → ACROSS THE PERIOD ELECTRON AFFINITY INCREASES greater effective nuclear charge, easily forms anions

70. Other Trends Reactivity of Metals Video 1 Reactivity of Metals Video 2

71. Increasing reactivity METAL REACTIVITY

72. Increasing reactivity NONMETAL REACTIVITY