1. Ch. 8 Covalent Bonding 8.1 Molecular Compounds

2. I. Molecules A. Neutral groups of atoms joined by covalent bonds B. Covalent bonds: atoms share electrons instead of giving/taking when electronegativity similar Ex. CH 4, CO 2, H 2 O, NH 3 C. Diatomic molecule: pairs of same atoms (O 2 )

3. II. Covalent Properties A. Weaker than ionic bonds B. Lower melting/boiling pts. C. Between non-metals D. Usually gas or liquid at room temp.

4. 8.2 The Nature of Covalent Bonding

5. I. Single Covalent Bond A. Sharing of a single electron pair between atoms B. Can represent e - pair as line C. Unshared e - s called “lone pair”, e - s between atoms represent the bond

6. II. Octet Rule A. Still follow rule that atoms are most stable with 8 valence electrons B. Octet comes from atoms own e - and shared ones

7. III. Multiple Bonds A. Double covalent bond: sharing two PAIRS of electrons B. Triple covalent bond: sharing three PAIRS of electrons Ex. N 2

8. IV. Lewis Dot Models A. Count up valence electrons of all atoms B. Usually first atom in formula is in center C. Start with single bonds (two e - ), see if remaining e - can be used to make all stable D. Try double or triple bonds if not E. Molecule charges mean missing or extra collective electrons (use brackets around molecule)

9. V. Coordinate Covalent Bonds A. When one atom contributes both electrons for a bond B. Ex.: Carbon Monoxide (CO) If electrons shared equally: Since not stable, you get:

10. VI. Resonance Structures A. Have two or more equally stable electron dot structures B. Rotate between structures by e - moving Example: Ozone

11. VII. Octet Exceptions A. The octet rule cannot be met in molecules with odd numbers of e - B. Some atoms can have more or less than 8 valence e - and be stable Boron trifluoride

12. VIII. Formal Charge A. If two different e - dot structures can both be stable, formal charge can determine which is more stable B. Formal charge = valence e - - e - in molecule for atom C. E - for atom = lone pairs + ½ of bonding e - D. More stable = less charges and any negatives on most electronegative atoms

13. IX. Bond Dissociation Energies A. Energy needed to break apart two covalently bonded atoms B. Stronger for multiple bonds Ex. C – C = 347 kiloJoule (kJ) C = C = 614 kJ C  C = 839 kJ C. Reaction energy = sum of energy to break bonds (reactants) – sum of energy to make bonds (products)

14. 8.3 Bonding Theories

15. I. Hybrid Orbitals A. Combination of orbitals of an atom B. Names based on orbitals and # of e - involved C. When 4 atoms/lone pairs attached to central atom there is one s and three p orbitals used (sp 3 )

16. II. Double/Triple Bond Hybrids A. When 3 atoms/lone pairs attached to central atom use one s and two p (sp 2 ) B. Left-over P orbital perpendicular to rest C. 2 atoms/lone pairs use 1 s and 1 p (sp), two p perpendicular

17. III. More Than an Octet A. To exceed 8 valence electrons, atoms use d orbitals B. 10 e - (sp 3 d) C. 12 e - (sp 3 d 2 ) D. Double and triple bonds need the same hybrid orbitals as single bonds and lone pairs

18. IV. Molecular Orbitals A. E - fit in orbitals in atoms, when molecules combine molecules have own orbitals B. Two S orbitals overlap to form sigma (σ) orbitals with σ bonds

19. C. P-orbitals can align or be parallel to each other D. Aligning form σ orbitals, parallel form Pi (  ) orbitals E. σ bonds stronger than  bonds F. Single covalent bonds are made of σ bonds G. Double bonds = 1 σ and 1 , triple = 1 σ and 2 

20. Hybrid orbitals attach to form molecular orbitals Ex. C 2 H 4 Ex. C 2 H 2

21. V. VSEPR Theory A. Shows 3-D structure of molecules B. Valence-Shell Electron-Pair Repulsion: because e - repel each other, 3-D shape puts pairs farthest apart C. Use shape that puts E - PAIRS and BONDED ATOMS farthest apart

22. VI. Examples A. 4 bonds = tetrahedral B. 3 bonds, 1 lone pair = pyramidal

23. C. 2 bonds, 2 lone pairs = bent D. 2 bonds = linear (CO 2 )

24. VII. Other VSEPR Shapes Square Pyramidal See-Saw

25. 8.4 Polar Bonds and Molecules

26. I. Bond Determination B. Difference 2 or more = ionic bond C. 0.4 or less difference = covalent bond D. 0.4 - 2 difference = polar covalent bond A. Difference between ionic and covalent bond is electronegativity of atoms

27. II. Polar Bond A. Unequal sharing of electrons B. Cl (3) pulls e - closer to it than H (2.1), Cl gets slight negative charge (  - ), H slight positive (  + )

28. III. Polar Molecule A. When polar bond makes entire molecule polar (“dipole”) B. Sometimes polar bonds don’t make molecule polar

29. IV. Intermolecular Attractions A. Molecules attracting each other B. Dipole attractions: charge of one molecule attracted to opposite charge on another Ex. Dipole-Dipole

30. V. Induced dipoles When electrons temporarily shift positions to form artificial dipoles (“Van der waals Forces”/ “London Dispersion Forces”)

31. VI. Hydrogen Bonds Hydrogen in a polar molecule attracted to electronegative atom on neighboring molecule Example: H 2 O Hydrogen Bond

32. VII. Network Solids A. Most covalent molecules melt at low temp. B. Network solids are very stable covalently bonded molecules with high melting pts. Vaporizes at 3500ºC Silicon Carbide melts at 2700ºC