1. Outline of Topics 1. Atomic Radius 2. Electronegativity 3. Ionization Energy 4. Ions 5. Ionic Compounds 6. Covalent Compounds 7. Lewis Structures

2. Trendy Table OBJECTIVE: Learn 3 patterns of the Periodic Table

3. Trendy Table 3 trends/patterns in the PT 1.Atomic Radius 2.Ionization Energy 3.Electronegativity

4. Trendy Table Atomic Radius Ionization Energy Electronegativity Definition Period INCREASES Row INCREASES

5. 1. Atomic Radius 1 st Trend - Atomic Radius What is it? Size of Atom

6. 1. Atomic Radius 1 st Trend - Atomic Radius As you go down a group, the number of electrons… So as you go down a group, the size of the atom…

7. 1. Atomic Radius 1s 1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1

8. 1. Atomic Radius What do you notice going across periods?

9. 1. Atomic Radius Atomic Radii – Size of Atom So, as you go down a period, atomic radius ______. As you go across a period…wait WHY?

10. 1. Atomic Radius 1s 1 1s 2 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 1

11. 1. Atomic Radius Atomic Radii – Size of Atom Nuclear Charge

12. 1. Atomic Radius NO ENERGY LEVEL added across periods

13. 1. Atomic Radius Atomic Radius – Size of Atom

14. 1. Atomic Radius 3 Patterns 1.Atomic Radius 2.Electronegativity 3.Ionization Energy Atomic Radius ElectronegativityIonization Energy Period INCREASES Row INCREASES

15. 1. Trendy Table Periodic Trends Atomic Radii – Size of Atom

16. 1. Trendy Table Periodic Trends Atomic Radii – Size of Atom

17. 1. Atomic Radius Summary & Review 1.What is atomic radius? 2.What happens to radius going down AND WHY? 3.What happens to radius across from left to right and WHY? 4.Which have smallest and largest radius? 5.Arrange by radius from smallest to largest carbon, cesium, copper, helium, iron, potassium

18. 2. Electronegativity 2 nd Trend: Electronegativity What is this? electronegativity = wanting/attracting electrons

19. 2. Electronegativity 2 nd Trend: Electronegativity related to atomic radius smaller the radius, higher the electronegativity higher electro = really want/good @ attract e-

20. 2. Electronegativity 3 Patterns 1.Atomic Radius 2.Electronegativity Atomic Radius ElectronegativityIonization Energy Period INCREASES Row INCREASES

21. 2. ELECTRONEGATIVITY 2 nd Trend: Electronegativity

22. 2. Electronegativity 2 rd Trend: Electronegativity

23. 2. Electronegativity 2 rd Trend: Electronegativity Fluorine has highest electronegativity Francium has lowest electronegativity Noble Gases do not count – why?

24. 3 rd Trend: Electronegativity

25. 2. Electronegativity 3 rd Trend: Electronegativity

26. 2. Electronegativity Summary & Review 1.What is electronegativity? 2.Which group is on the 0 of the y-axis? 3.Which element has highest electronegativity, and which has lowest? 4.Arrange from lowest to greatest electronegativity oxygen, cesium, sulfur, potassium, zinc, fluorine

27. 3. Ionization Energy 3 rd Trend: Ionization Energy What is this? energy needed to remove/steal an electron from an atom

28. 3. Ionization Energy 3 nd Trend: Ionization Energy What is this? Energy needed to remove an electron from an atom OPPOSITE of ATOMIC RADIUS

29. 3. Ionization Energy 3 Patterns 1.Atomic Radius 2.Electronegativity 3.Ionization Energy Atomic Radius ElectronegativityIonization Energy Period INCREASES Row INCREASES

30. 3. Ionization Energy 3 nd Trend: Ionization Energy

31. 3. Trendy Table How to tell the difference between the two? NOBLE GASES do not care about Electronegativity

32. 3. Ionization Energy 3 nd Trend: Ionization Energy

33. Periodic Trends 3 Patterns 1.Atomic Radius – size of atom 2.Ionization Energy – energy to remove electrons 3.Electronegativity – ability to attract electrons Atomic Radius Ionization Energy Electronegativity Period INCREASES Smallest: He Largest: Cs INCREASES Weakest: Cs Strongest: He INCREASES Weakest: Cs Strongest: F Row INCREASES Smallest: He Largest: Cs INCREASES Weakest: Cs Strongest: He INCREASES Weakest: Cs Strongest: F

34. 3. Trendy Table

35. 3. Ionization Energy Summary & Review 1.What is ionization energy? 2.What are the relationships between atomic radius, electronegativity, and ionization energy? 3.Arrange from lowest to greatest ionization energy copper, neon, silicon, cesium, helium, phosphorus, calcium, fluorine,

36. 3. Trendy Table Periodic Trends 3 nd Trend: Ionization Energy What is this? Energy needed to remove an electron from an atom A filled orbital is a happy and stable orbital All elements want to be NOBLE

39. 4. IONS 5 + - 2 = 5 + - 3 = 5 + - 4 = 5 + - 6 =

40. 4. IONS OBJECTIVE: When atoms give away electrons

41. 4. IONS ION What is it? An atom that LOST/GAINED electron(s)

42. 4. IONS How do we know if an atom GAINS or LOSES electrons? and How MANY electrons is it going to gain or lose?

43. 4. IONS Goal is to get to the NEAREST Noble Gas RULES 1.If you go LEFT, you become POSITIVE. 2.If you go RIGHT, you become NEGATIVE

44. 4. IONS Ions that LOSE electrons have a POSTIVE charge called cations pronounced “cat – ions”

45. 4. IONS All metals become CATIONS

46. 4. IONS Transition Metals do not care about being NOBLE They have their own thing going one. DO NOT HAVE TO MEMORIZE…kind of

47. 4. IONS Ions that GAINS electrons have a NEGATIVE charge called anions pronounced “an – ions”

48. 4. IONS Ions with at NEGATIVE charge called anions pronounced “an – ions”

49. 4. IONS Helium Neon Argon Krypton 1.Write the electron configuration for the following elements 2.UNDERLINE valence electrons The S orbital can hold how many electrons MAX? TWO The P orbital can hold how many electrons MAX? SIX Noble Gases do not become ions because their orbitals are FILLED THEY have all the VALENCE ELECTRONS THEY NEED

50. 4. IONS Helium 1s 2 Neon [He]2s 2 2p 6 Argon [Ne]3s 2 3p 6 Krypton [Ar]4s 2 3d 10 4p 6 Xenon [Kr]5s 2 4d 10 5p 6 Alkali and Alkali Earth Metals want to be noble by having the SAME electron configuration as the nearest noble gas Non-Metals want to be noble by having the SAME electron configuration as the nearest noble gas

51. 4. IONS So far we learned how a “normal” atom becomes an ion Li  Li + + 1electron Mg  Mg 2+ + 2electrons N + 3electrons  N 3- O + 2electrons  O 2- F + electron  F -

52. 4. IONS Ex: CN - NH 4 + CO 3 2- Because these ions are made from two or more atoms, they are called polyatomic ions Some ions are made from two or more atoms

53. 4. IONS Because these ions are made from two or more atoms, they are called polyatomic ions -ite has less oxygen than –ate Memorize? Kind of

54. 4. IONS Summary & Review 1.What is an ion? 2.What are the two types of ions? 3.Metals form these types of ions, and non-metals form these types of ions. 4.Why do Noble Gases not become ions? 5.Predict the charge of each atom if it were to become an ion: Ca, Cl, Cu, F, Fe, Li, Na, N, O

55. 5. Ionic Compounds 4x + -16 = 0 2x + -8 = 0 3x + -12 = 0 2 + -x = 0 10 + -2x = 0 12 + -3x = 0

56. 5. Ionic Compounds OBJECTIVE: Joining Cations + Anions to make compounds

57. 5. Ionic Compounds Na + + Cl -  NaCl cation anionionic compound forumula NaCl joined by IONIC BOND IONIC BOND = giving away/gaining electrons

58. 5. Ionic Compounds Li + + Cl -  K + + Cl -  Ca 2+ + Cl -  Li + + O 2-  Ca 2+ + O 2-  LiCl KCl CaCl 2 Li 2 O CaO Charges must cancel and = 0

59. 1. Separate into cations and anions 2. Separate into Alkali, Alkali Earth, Halogens, Transition Metals, and Polyatomic Ions 3. Make the following compounds 5. Ionic Compounds H + + Cl -  Na + + Cl -  Mg 2+ + Cl -  Mg 2+ + SO 4 2-  Mg 2+ + SO 3 2-  Fe 2+ + Cr 2 O 7 2-  NH 4 + + Cl -  Fe 2+ + O 2-  Na + + O 2-  Na + + OH -  Fe 3+ + OH -  NH 4 + + SO 4 2- 

60. 5. Ionic Compounds Cu + + CH 3 COO -  NH 4 + + Cr 2 O 7 -  Mg 2+ + OH -  CuCH 3 COO NH 4 Cr 2 O 7 Mg ( OH ) 2 Parenthesis for POLYATOMIC

61. 6. Naming Ionic Compounds Write formula for ionic compounds 1. H + + SO 4 2-  2. Al 3+ + O 2-  3. Mg 2+ + OH -  4. Fe 2+ + Cl -  5. Na + + HCO 3 -  6. Ca 2+ + PO 4 3-  7. Na + + NO 2 -  8. Fe 3+ + Cl -  9. NH 4 + + PO 4 3- 

62. Summary & Review What is the correct formula for magnesium hydroxide: MgOH 2, MgO 2 H 2 or Mg(OH) 2 ? K+K+ Fe 3+ NH 4 + Ba 2+ Cl - KCl SO 4 2- PO 4 3- NO 3 - OH -

63. Complete table on separate paper What is the correct formula for magnesium hydroxide: MgOH 2, MgO 2 H 2 or Mg(OH) 2 ? K+K+ Fe 3+ NH 4 + Ba 2+ Cl - KClFeCl 3 NH 4 ClBaCl 2 SO 4 2- K 2 SO 4 Fe 2 (SO 4 ) 3 (NH 4 ) 2 SO 4 BaSO 4 PO 4 3- K 3 PO 4 FePO 4 (NH 4 ) 3 PO 4 Ba 3 (PO 4 ) 2 NO 3 - KNO 3 Fe(NO 3 ) 3 NH 4 NO 3 Ba(NO 3 ) 2 OH - KOH Fe(OH) 3 NH 4 OH Ba(OH) 2

64. Complete table on separte paper ? mols ? grams K+K+ Fe 3+ NH 4 + Ba 2+ Cl - 12g 3.27 mol SO 4 2- 79.3g 65.30 mol PO 4 3- 123.4g 2.31 mol NO 3 - 423.1g 3.27 mol OH - 1.497g 1.985 mol

65. 7. Lewis Structures How many valence electrons do the each of the Noble Gases have? He Ne Ar Kr Xe Rn

66. 7. Lewis Structures OBJECTIVE: How bonds are made

67. 7. Lewis Structures Most elements want 8 valence electrons Octet Rule

68. 7. Lewis Structures dot = valence electron

69. 7. Lewis Structures blue line = bond

70. 7. Lewis Structures Elements that do not want 8 dots Group 1 H, Li, Na, K, etc. Only wants 2 dots more exceptions later

71. 7. Lewis Structures Draw the following Lewis Structures H 2 Cl 2 H 2 SHBr OF 2 H 2 O NH 3 CH 4 CH 2 F 2 C 2 H 6

72. 6. Covalent Bonds & Compounds Most compounds = covalent bonds

73. 7. Lewis Structures Draw the following structures O2N2O2N2 COCO 2 HCNSO 3 C 2 Cl 4 COCl 2

74. 7. Lewis Structures 1. Write elements 2. Count total number of valence e 3. LEAST electronegative is in the center (usually CARBON) 4. Draw bond 5. Then draw dots until you run out 6. Make double/triple bonds if necessary

75. 7. Lewis Structures Summary & Review 1. What element goes in the middle of a Lewis Structure? 2. Which group does not need 8 dots? 3. Which element is NEVER in the center

76. 4. Lewis Dot Structures

77. What are the two poles of a magnet called? What do these poles mean/tell you? When do you use this diagram in your other classes?

78. 6. Covalent Bonds & Compounds OBJECTIVE: When atoms SHARE electrons

79. 6. Covalent Bonds & Compounds co-valent So which electrons are the atoms sharing?? sharing VALENCE electrons

80. Difference between Ionic & Covalent IONIC BOND IONIC COMPOUND Covalent Bond Covalent Compound Atoms GAIN or LOSE electrons Atoms SHARE electrons Na + + Cl -  NaCl 6. Covalent Bonds & Compounds C + Cl  CCl 4

81. 6. Covalent Bonds & Compounds Covalent Bond/Compounds = when atoms SHARE valence electrons video

82. 6. Covalent Bonds & Compounds Some examples of covalent compounds

83. 6. Covalent Bonds & Compounds Some example of covalent compounds

84. 6. Covalent Bonds 2 types Nonpolar Covalent valence electrons EQUALLY shared Polar Covalent valence electrons NOT equally shared

85. 6. Covalent Bonds Electronegativity determines type 0 – 0.49 Nonpolar C 0.50 – 2.0 Polar C above 2.0 Ionic

86. 6. Covalent Bonds H 2, Cl 2, F 2, N 2, NO 1. Draw Lewis Structure 2. Subtract electronegativity values to determine bond type. 3. All have values less than 0.49, so all the above share electrons equally

87. 6. Covalent Bonds HBr, NaCl, OF 2, NH 3, KCl 1. Draw Lewis Structure 2. Subtract electronegativity values to determine bond type

88. 6. Covalent Bonds HBr, NaCl, OF 2, NH 3, KCl What does it mean for HBr, OF2, and NH3 to be polar? Means shared electron spends more time with… the element with higher electronegativity

89. 6. Covalent Bonds If polar you use the signs d- and d+ What do these signs mean??? Means shared electron spends more time on the d- side, so d- side has a small negative charge So how do you know which side is d-? the element with higher electronegativity

90. 6. Covalent Bonds If Ionic, we just use - and + What do these signs mean??? Minus side has a negative charge because it gained and electron So how do you know which side is -? anion is the minus side

91. 3. Types of Covalent Bonds Difference > 0.5 Difference > 2.1 Polar Ionic Least Ionic Most Ionic

92. 6. Covalent Bonds & Compounds Summary & Review 1. How is an ionic bond/compound different from a covalent bond/compound? 2. What is shared in a covalent bond/compound? 3. How can you tell if bond is nonpolar, polar, or ionic? 4. If a bond is polar, what are the two signs you use? 5. How do you know where put the signs? 6. What do those signs mean?