Chemical Bonding

1. Chemical Bonding

2. Chemical compounds A. Atoms Combine 1. To become Stable • Atoms are stable when they have a full valence energy level • Octet rule: Usually 8 e- • Exception: He is stable w/ 2 e- • Atoms gain, lose or share e- to achieve a full valence level c. Atoms are held together by a chemical bond

3. Chemical Bonds, Lewis Symbols, and the Octet Rule • Chemical bond: attractive force holding two or more atoms together. • Ionic bondresults from the transfer of electrons from a metal to a nonmetal. • Covalent bondresults from sharing electrons between the atoms. Usually found between nonmetals. • Metallic bondattractive force holding pure metals together.

4. Ionic Compounds • Ionic Bonds 1.Are formed a.between (+) and (-) ions b.Whenone atom gains e- while the other loses e- 2.Compounds with ionic bonds are called ionic compounds

5. Ionic Bonds: One Big Greedy Thief Dog!

6. Covalent Bonding B. Covalent Bonds 1.Are formed when atoms share e- 2.Even sharing results in non-polar covalent bonds 2.Uneven sharing results in polar covalent bonds

7. Polar Covalent Bonds: Unevenly matched, but willing to share.

8. Metallic Bonding C. Formed between atoms of metallic elements 1. Electron cloud around atoms 2. Good conductors at all states, lustrous, very high melting points Examples; Na, Fe, Al, Au, Co

9. Metallic Bonds: Mellow dogs with plenty of bones to go around.

10. Covalent Bonding B. Covalent Bonds 2.Are formed Between two nonmetals 3. Compounds with covalent bonds are called molecules

11. Chemical Bonds, Lewis Symbols, and the Octet Rule Lewis Symbols

12. Learning Check  A. X would be the electron dot formula for 1) Na 2) K 3) Al   B.  X  would be the electron dot formula  1) B 2) N 3) P

13. Drawing Lewis Structures • Follow Step by Step Method • Total all valence electrons. [Consider Charge] • Write symbols for the atoms and guess skeleton structure [ define a central atom ]. • Place a pair of electrons in each bond. • Complete octets of surrounding atoms. [ H = 2 only ] • Place leftover electrons in pairs on the central atom. • If there are not enough electrons to give the central atom an octet, look for multiple bonds by transferring electrons until each atom has eight electrons around it.

14. Exceptions to the Octet Rule • Central Atoms Having Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3, with only 6 • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

15. Molecular Shapes: VSEPR • There are five fundamental geometries for molecular shape:

16. Linear • 2 atoms attached to center atom • 0 unshared pairs (lone pairs) • Bond angle = 180o • Type: AB2 • Ex. : BeF2

17. Trigonal Planar • 3 atoms attached to center atom • 0 lone pairs • Bond angle = 120o • Type: AB3 • Ex. : AlF3

18. Tetrahedral • 4 atoms attached to center atom • 0 lone pairs • Bond angle = 109.5o • Type: AB4 • Ex. : CH4

19. Trigonal Bipyramidal • 5 atoms attached to center atom • 0 lone pairs • Bond angle = • equatorial -> 120o • axial -> 90o • Type: AB5 • Ex. : PF5

20. Octahedral • 6 atoms attached to center atom • 0 lone pairs • Bond angle = 90o • Type: AB6 • Ex. : SF6

21. These are for molecules with both paired and unshared (lone) pairs of electrons around the central atom.

22. Bent • 2 atoms attached to center atom • 2 lone pairs • Bond angle = 104.5o • Type: AB2E2 • Ex. : H2O

23. Trigonal Pyramidal • 3 atoms attached to center atom • 1 lone pair • Bond angle = 107o • Type: AB3E • Ex. : NH3

25. Figure 9.3 HyperChem

26. Summary of VSEPR Molecular Shapes

27. Learning Check Determine the molecular geometry of each of the following: A. CCl4 1) tetrahedral 2) pyramidal 3) angular B. SO3 1) trigonal planar 2) pyramidal 3) angular C. PCl3 1) trigonal planar 2) pyramidal 3) angular Timberlake LecturePLUS

28. Solution Determine the molecular geometry of each of the following: A. CCl4 1) tetrahedral B. SO3 1) trigonal planar C. PCl3 2) pyramidal Timberlake LecturePLUS

30. Polarity

31. O O S Just as electrons push away from each other, so do molecules HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: London dispersion forces, caused by “temporary dipoles”. SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

32. O O S Intermolecular Forces What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: dispersion forces. SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

33. or … … H H B A A A Intermolecular Forces Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A & B are N, O, or F

34. Intermolecular Forces Dispersion Forces Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules ion-induced dipole interaction dipole-induced dipole interaction