1. Review Inorganic Chemistry Chemistry 20

2. More than 20 million compounds are composed of these 116 elements. Element: is a substance consists of identical atoms. Cannot be divided by chemical & physical methods. Carbon, Hydrogen, Oxygen 116 elements – 88 in nature

4. main-group elements: 1A to 8A transition elements: 1B to 8B (3 – 12) inner transition elements: between B3 & B4 (58 to 71 and 90 to 103) Column: same properties (main group) row or period

5. nonmetals metals metalloids Classification of the elements metals: solid (except mercury), shiny, conductors of electricity, ductile, malleable nonmetals: solid, liquid or gas, do not conduct electricity (except graphite) metalloids: between metals and nonmetals t o

6. n=1 n=2 n=3 n=4 n=5 E1E1 E2E2 E3E3 E4E4 E5E5 Bohr model Principal energy levels or Shell 324 183 82 21 Maximum number of electrons Level Lower energy Higher energy

7. Sublevel:spdf Orbital: is a region of space and it can hold 2 electrons (max). s p d f 2 2+2+2=6 2+2+2+2+2=102+2+2+2+2+2+2=14

8. LevelOrbitals Maximum number of electrons 11s2 22s, 2p2 + 6 = 8 33s, 3p, 3d2 + 6 + 10 = 18 44s, 4p, 4d, 4f2 + 6 + 10 +14 = 32 3 2 3d 3p 3s 2p 2s Principal energy levelOrbitals Energy Order of filling 1 1s

9. Orbital box diagrams H (1) 1s 1s 1 Electron configuration 1s He (2) 1s 2 Li (3) 1s 2s 1s 2 2s 1 C (6) 1s 2s 1s 2 2s 2 2p 2 2p x 2p y 2p z Electrons configuration: description of the orbitals that its electrons occupy.

10. Valence electrons: outer-level electrons Valence level: outer level Cl (17) 1s 2 2s 2 2p 6 3s 2 3p 5 7 valence electrons Ar (18) 1s 2 2s 2 2p 6 3s 2 3p 6 8 valence electrons C (6) 1s 2 2s 2 2p 2 4 valence electrons Ne (10) 1s 2 2s 2 2p 6 8 valence electrons Noble gases Filled valence levels

11. Lewis dot structure H He Li C Cl 1A 2A 3A 4A 5A 6A 7A 8A

12. Level 1 main-group elements 1A – 8A Maximum 2 electrons in valence level Hydrogen and Helium Other Levels Maximum 8 electrons in valence level

13. 1.Ionic bonds 2. Covalent bonds Chemical Bonds

14. Goal of atoms Filled valence level Noble gases (Stable) Na + : 1s 2 2s 2 2p 6 + e - Na: 1s 2 2s 2 2p 6 3s 1 Ne: 1s 2 2s 2 2p 6 Ar: 1s 2 2s 2 2p 6 3s 2 3p 6 Octet rule

15. Goal of atoms Filled valence level + e - Ar: 1s 2 2s 2 2p 6 3s 2 3p 6 Noble gases (Stable) Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 Cl - : 1s 2 2s 2 2p 6 3s 2 3p 6 Na + : 1s 2 2s 2 2p 6 + e - Na: 1s 2 2s 2 2p 6 3s 1 Ne: 1s 2 2s 2 2p 6

16. Ionic bonds Metal-Nonmetal Na: 1s 2 2s 2 2p 6 3s 1 Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 Anion Cation Na + : 1s 2 2s 2 2p 6 Cl - : 1s 2 2s 2 2p 6 3s 2 3p 6

17. Covalent bonds Nonmetal-Nonmetal Metalloid-Nonmetal Sharing of valence electrons

18. Lewis Dot Structure H He Li C Al NCl H H Or H H Cl H Lewis Structure H Cl Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 H: 1s 1 He: 1s 2 Ar: 1s 2 2s 2 2p 6 3s 2 3p 6

19. Covalent compounds H – C – H – O CH 2 O H C H O H – C – H – O H C H O – Correct H C H H H H – C – H – H – H CH 4 H – N – H – H NH 3 H N H H

20. H C CH C – C C2H4C2H4 H H H H H H H C CH C = C H H H H H H H C CH H – C – C – H C2H2C2H2 H C CH H – C C – H Correct

21. Electronegativity A measure of an atom’s attraction for the electrons

22. Covalent bonds Nonpolar covalent bond:electrons are shared equally. Polar covalent bond:electrons are shared unequally. HCl δ+ δ- Dipole

23. Polarity 1.Molecule has polar bonds. 2.Its centers of δ+ and δ- lie at different places (sides). O = C = O δ- δ+ δ- nonpolar molecule H – C – H – H – H δ+ δ- C = O H H δ+ δ- polar molecule N H H H δ- δ+

24. Intermolecular Forces London dispersion forces Dipole-dipole interaction Hydrogen bonding Ionic bonds Covalent bonds < Intramolecular (Bonding) Forces Intermolecular Forces

25. London dispersion forces Attractive forces between all molecules Only forces between nonpolar covalent molecules 2+ No Polarity Original Temporary Dipole δ-δ+ + 2+ He Original Temporary Dipole Induced Temporary Dipole __ _ _ He 2+ __ _ _ δ-δ+ He _ _ 2+ δ-δ+ He _ _ 2+

26. Dipole-Dipole Interactions Attractive force between two polar molecules stronger than London dispersion forces boiling point ↑

27. Hydrogen Bonds stronger than Dipole-Dipole and London dispersion forces surface tension High boiling point H2OH2O Between H bonded to O, N, or F (high electronegativity) → δ+ and a nearby O, N, or F → δ-

28. Hydrogen bonding CH 3 COOH Acetic acid δ- δ+

29. Factors that affect boiling point: 1. Intermolecular forces: London dispersion forces < Dipole-Dipole interactions < Hydrogen bonds 3. Molecular shape: With the same molecular weight. linear CH 3 -CH 2 -CH 2 -CH 2 -CH 3 > spherical CH 3 -C- CH 3 _ _ CH 3 2. Number of sites for intermolecular interaction (surface area): Larger surface areas (more electrons)  more sites for L.D.F   b.p. CH 3 -CH 2 -CH 2 -CH 2 -CH 3 > CH 3 -CH 2 -CH 3 Boiling point