2. Naming with Roman numerals Naming with Greek Prefixes

3. The Covalent Bond Is due to sharing (co-) valence electrons (- valent).

4. Covalent Bonds Form Hybridized Electron Orbitals

5. The Single Covalent Bond Atoms may share 1 electron pair = single covalent bond. Water methane Hydrogen gas

6. The Double Covalent Bond Atoms may share 2 electron pairs = double covalent bond. Oxygen gas

7. The Triple Covalent Bond Atoms may share 3 electron pairs = triple covalent bond. http://bcs.whfreeman.com/thelifewire/conten t/chp02/02020.html

8. Resonance Structures Alternating or average possible positions of a double covalent bond in a compound.

9. Coordinate-Covalent Bond When one shared electron pair comes entirely from one atom. NH 3 + HCl  NH 4 + + Cl -

10. Bond Length Bond length= the distance between the nuclei. Bond Enthalpy = the energy needed to break the bond and separate the atoms.

11. Representing Covalent Compounds Using Lewis Dot Pictures Hint: One line = one shared pair of electrons. Dots on outside are unshared electrons = saturated = unsaturated Hint: The most electronegative atom is usually in the center of Lewis Dot pictures.

12. Remember: PrefixNumber Value Hemi-½ Mono-1 Di-2 Tri-3 Tetra-4 Penta-5 Hexa-6 Hepta-7 Octa-8 Nona-9 Deca-10 Doci-12

13. Covalent Nomenclature Naming covalent compounds is easy Name the first compound and the second ends in ‘–ide’ (just like ionic), but: Use number prefixes to tell how many. Examples: CCl 4 = N 2 O 3 = CO = carbon tetrachloride dinitrogen trioxide carbon monoxide Hint: Do not use ‘mono-’ on the first element.

14. You Try 1.CF 4 2.As 2 O 3 3.CO 2 4.SO 2 5.NF 3 6.P 2 O 5 Carbon tetrafluoride Diarsenic trioxide Carbon dioxide Nitrogen trifluoride Sulfur dioxide Diphosphorous pentoxide Hint: Drop the double vowel, so its not “pentaoxide”, its “pentoxide”.

15. Covalent Structure Covalent compounds have hybridized molecular orbitals, such as sp 3. VSEPR = Valence Shell Electron Pair Repulsion

16. Molecular Geometry Remember: Molecules are 3-dimensional structures.

17. You Try! Name the shape: Linear 1. 2. Trigonal pyramidal 3. Tetrahedral 4. Octahedron Trigonal bipyramidal 5.

18. Polar and Nonpolar Covalent Bonds Electrons may be shared very evenly = nonpolar covalent bond. Electrons may be shared unevenly = polar covalent bond. Hint: Lower case delta, δ, means “partial,” as in a partial charge(δ + or δ - ).

19. Polar and Nonpolar Bonds and Molecules H 2 has a nonpolar bond in nonpolar molecule because the pulls are equal Cl 2 has a nonpolar bond in nonpolar molecule because the pulls are equal CCl 4 has polar bonds in a nonpolar molecule because the pulls cancel HBr and H 2 O have polar bonds in polar molecules because the pulls are unequal and do not cancel

20. Using Electronegativity to Determine Bond Type and Character Electronegativity is a number assigned to atoms to help determine the pull on electrons and the type of bond that forms. The most commonly used method of calculation is that originally proposed by Linus Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from around 0.7 to 3.98 (hydrogen = 2.20). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in Pauling units. - WikipediaLinus Pauling dimensionless quantityhydrogen

21. Electronegativity difference: 0------------------------------------0.4-------------------------------2.1--------------------------3.3 Determining Bond Polarity and Type (Usually) H2H2 HF NaCl nonpolar covalent

22. Na 0.9 Cl 3.2 a) 3.0-3.0 = 0.0 = nonpolar b) 3.0-2.1 = 0.9 = polar 0------------------------------------0.4-------------------------------2.1--------------------------3.3 -----------nonpolar---------------------------polar----------------------------ionic--------------- c) 3.2-0.9 = 2.3 = ionic abc ionic bond covalent You Try: Show the math

23. You Decide: A) Non polar covalent bond B) Polar covalent bond C) ionic bond 1,2. HCl 3,4. CO 2 5,6. LiF 7,8. SO 2 9,10. O 2 11,12. KBr 13, 14. BH 3 15,16. H 2 O B) Polar 3.2 – 2.2 =1.0 B) Polar 3.4 – 2.6 =0.8 C) Ionic 4.0 – 1.0 = 3.0 B) Polar 3.4 – 2.6 = 0.8 A) Nonpolar 3.4 – 3.4 = 0.0 C) Ionic 3.0 – 0.8 = 2.2 B) Polar 3.4 – 2.2 = 1.2 A) Nonpolar 2.2 – 2.0= 0.2 How about the compound? polar nonpolar Not a molecule (ionic)! polar nonpolar Not a molecule (ionic)! nonpolar Polar and now the rest of the story…

24. Van der Waals Forces Van der Waals forces are intermolecular attractions between atoms caused by permanent or temporary dipoles. These are the two types you need to know… London Dispersion Forces and Hydrogen bonds.

25. London Forces London Dispersion forces - As electrons move around atoms, charge may briefly accumulate on one side of the atom.

26. Hydrogen Bonding The ends of polar molecules attract each other and can form Hydrogen Bonds. Hydrogen Bonds are intermolecular not atomic and are very important in nature.

27. Nonpolar substance tend to vaporize easily. Water’s properties depend on hydrogen bonds: Hydrogen bonds make it possible to have liquid, ice, and water vapor all at once on earth. Hydrogen bonds make water wet (give water cohesion and adhesion). Hydrogen bonds are the reason ice floats. Hydrogen bonds hold DNA together!! Hydrogen Bonds make life possible!! Do you recognize this molecule? Covalent Bond Hydrogen Bond

28. Water, “The Universal Solvent”: Hydrogen Bonds in Water pull apart salts = dissolve Water is a good solvent of polar and ionic compounds. Hint: Like dissolves like. + Hydrogen bonds make water “the universal solvent “.

29. The End (for now…)