Covalent Compounds

1. Covalent Compounds

2. Naming with Roman numerals Naming with Greek Prefixes

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

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

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

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

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

8. Coordinate-Covalent Bond • When one shared electron pair comes entirely from one atom. NH3 + HCl NH4+ + Cl-

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

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

11. Remember: Prefix Number 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

12. 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: CCl4 = N2O3 = CO = carbon tetrachloride Hint: Do not use ‘mono-’ on the first element. dinitrogen trioxide carbon monoxide

13. You Try Carbon tetrafluoride • CF4 • As2O3 • CO2 • SO2 • NF3 • P2O5 Hint: Drop the double vowel, so its not “pentaoxide”, its “pentoxide”. Diarsenic trioxide Carbon dioxide Sulfur dioxide Nitrogen trifluoride Diphosphorouspentoxide

14. Covalent Structure • Covalent compounds have hybridized molecular orbitals, such as sp3. VSEPR = Valence Shell Electron Pair Repulsion

15. Molecular Geometry • Remember: Molecules are 3-dimensional structures.

16. You Try! • Name the shape: 2. 1. Linear Trigonal pyramidal 3. 5. 4. Tetrahedral Trigonalbipyramidal Octahedron

17. 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-).

18. Polar and Nonpolar Bonds and Molecules H2 has a nonpolar bond in nonpolar molecule because the pulls are equal Cl2 has a nonpolar bond in nonpolar molecule because the pulls are equal CCl4 has polar bonds in a nonpolar molecule because the pulls cancel HBr and H2O have polar bonds in polar molecules because the pulls are unequal and do not cancel

19. 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. - Wikipedia

20. Determining Bond Polarity and Type (Usually) Electronegativity difference: 0------------------------------------0.4-------------------------------2.1--------------------------3.3 nonpolar covalent H2 NaCl HF

21. 0------------------------------------0.4-------------------------------2.1--------------------------3.30------------------------------------0.4-------------------------------2.1--------------------------3.3 -----------nonpolar---------------------------polar----------------------------ionic--------------- a b c covalent covalent covalent a) 3.0-3.0 = 0.0 = nonpolar Cl 3.2 Na 0.9 ionic bond covalent c) 3.2-0.9 = 2.3 = ionic b) 3.0-2.1 = 1.1 = polar

22. You Decide: A) Non polar covalent bond B) Polar covalent bondC) ionic bond 1,2. HCl 3,4. CO2 5,6. LiF 7,8. SO2 9,10. O2 11,12. KBr 13, 14. BH3 15,16. H2O How about the compound? B) Polar 3.2 – 2.2 =1.0 polar B) Polar 3.4 – 2.6 =0.8 nonpolar C) Ionic 4.0 – 1.0 = 3.0 Not a molecule (ionic)! B) Polar 3.4 – 2.6 = 0.8 polar A) Nonpolar 3.4 – 3.4 = 0.0 nonpolar C) Ionic 3.0 – 0.8 = 2.2 Not a molecule (ionic)! A) Nonpolar 2.2 – 2.0= 0.2 nonpolar B) Polar 3.4 – 2.2 = 1.2 Polar and now the rest of the story…

23. 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.

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

25. 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.

26. Hydrogen Bonds make life possible!! • 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!! Covalent Bond Hydrogen Bond Do you recognize this molecule?

27. 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 “.

28. The End (for now…)