1. Covalent/Molecular Bonding Ch. 16

2. The Nature of Covalent Bonding 16-1 Skip pgs 444 - 451

4. Covalent Bonds Covalent (molecular) bond = the attraction of two atoms for a shared pair of electrons –Neither atom will have an ionic charge –Usually between 2 nonmetals (some involve metalloids)! Covalent compound = a compound whose atoms are held together by covalent bonds Molecule = an uncharged group of two or more atoms held together by covalent bonds

5. Single Covalent Bonds Single Covalent Bond = 2 atoms share one pair of electrons. –H 2, F 2, H 2 O Structural Formula = chemical formula that shows the arrangement of atoms. –H + H  H H  H – H Element Compound Structural dots dots formula

6. – F + F  F F  F - F – H + O  O H  O - H H H H F2:F2: H 2 O:

7. H CH 4  C + 4H  HC H H H * HC H H NH 3  N + 3H  HNH H HNHHNH H

8. Double and Triple Covalent Bonds Double Covalent bonds = bonds that involve 2 shared pairs of electrons. –O 2  O O Triple Covalent bonds = bonds that involve 3 shared pairs of electrons. –N 2  N N

9. O 2  O + O  O O  O O

10. N 2  N + N  N N  N N

11. CO 2  O + C + O  O C O  O = C = O Add to notes, above N 2

12. Bonding Theories 16-2 Skip pgs. 452–454 + 457-459

13. VSEPR Theory VSEPR Theory states that because electron pairs repel, molecular shape adjusts so that valence-electron pairs are as far apart as possible. –Ex: H 2 O bond is NOT linear! 2H + O  O H H

14. VSEPR Geometrics A = Central Atom, X = Attached Species, E = Extra Pair of e-’s on A Total # of Attached Species Species Type Molecular Geometry Example 2AX 2 LinearCO 2 4AX 4 AX 3 E AX 2 E 2 Tetrahedral Pyramidal Bent CH 4 NH 3 H 2 O

15. A A A A A A 107° 105° Triatomic 120° PyramidalBent

16. Linear Example

17. Tetrahedral Example

18. Pyramidal Example

19. Bent Example

20. Polar Bonds + Molecules 16-3 Part I

21. Bond Polarity Bonding pairs of electrons are pulled, as in a tug-of-war, between nuclei of atoms sharing electrons. If bonding pairs are shared equally it is a nonpolar covalent bond. –Atoms will have equal electronegativities (pg. 405) –Ex: N 2, O 2, H 2, Cl 2, CO 2 If bonding pairs are shared unequally it is a polar covalent bond. –Atoms have unequal electronegativity. –H 2 O, HCl, CO

22. Polar Molecules Polar molecule = one end of molecule is slightly negative and other end is slightly positive. Ex: HCl Electronegativity: H = 2.1, Cl = 3.0 Difference = 0.9 Ex: H 2 O H = 2.1, O = 3.5 Difference = 1.4

23. Electronegativity Differences + Bond Types Electronegativity Difference Type of BondExample 0.0 – 0.3Nonpolar CovalentH – H (0.0) 0.4 – 1.0Moderate Polar Covalent ∂+ ∂- H – Cl (0.9) 1.1 – 2.0Very Polar Covalent ∂+ ∂- H – F (1.9) > 2.0IonicNa + Cl - (2.1)

24. The polarity of a molecule depends on the shape + orientation of the bonds. –Ex: CO 2 polarity cancels out since it is linear = nonpolar molecule –Ex: H 2 O poles add up due to its bent shape = polar molecule

25. Polar: F H C H H H N H H OH Nonpolar: H HCH H O C OHCH

26. Comparing Ionic + Molecular Properties CharacteristicIonic CmpdCovalent/ Molecular Cmpd Representative UnitFormula UnitMolecule Bond FormationTransfer e-’sShare pairs e-’s Type of ElementMetal + nonmetal2 Nonmetal (possible metalloid) Physical StateSolidS, L, or G Melting PointHigh (>300°C)Low (< 300°C) Solubility in WaterHighHigh to Low Electrical Conductivity as aqueous soln GoodPoor to none

27. Attractions Between Molecules 16-3 Part II

28. Attractions between Molecules van der Waals forces = weakest attraction (ionic + covalent are stronger). –Three types: 1. Dispersion forces 2. Dipole interactions 3. Hydrogen bonds.

29. 1) Dispersion forces = weakest of all molecular interactions; caused by motion of electrons. Increases as the number of electrons increases. Halogen diatomic molecules (F 2, Cl 2, Br 2, I 2 ) Fluorine + Chlorine have weak dispersion forces (less electrons); thus are gases at STP. Bromine (more electrons) is a liquid at STP, and Iodine (most electrons) is a solid at STP.

30. 2) Dipole Interactions = occurs when polar molecules are attracted to one another. –The slightly negative region of a polar molecule is attracted to the slightly positive region of another polar molecule

31. –When placed in an electric field, dipole molecules become oriented with respect to (-) and (+) charge

32. 3) Hydrogen bonds = attractive forces in which a hydrogen covalently bonded to a very electronegative atoms is also weakly bonded to an unshared electron pair of another electronegative atom. -Strongest of intermolecular (van der Waals) forces

33. Network Solids Most molecules are easy to break; however, a few molecular solids are very stable. Network Solids = solids in which all atoms are covalently bonded to each other. –Solid does not “melt” until 1000°C or higher, in which it vaporizes without melting at all. –Ex: Diamond; made of carbon, each carbon bonded to 4 other carbons. Quartz (SiO 2 ) C – C – C – C

34. Diamond + Silicon carbide (SiC)