1. Chemical Bonding

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

3. Figure 8.3: Ionic Bonding

4. Figure 8.5: Covalent Bonding

5. Chemical Bonds Bond Type Single Double Triple # of e’s 2 4 6 Notation — =  Bond order 1 2 3 Bond strength Increases from Single to Triple Bond lengthDecreases from Single to Triple

7. Strengths of Covalent Bonds

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

9. The Octet Rule All noble gases except He has an s 2 p 6 configuration. Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). Caution: there are many exceptions to the octet rule. Chemical Bonds, Lewis Symbols, and the Octet Rule

10. Bond Polarity and Electronegativity Electronegativity Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Electronegativity increases across a period and down a group.

11. Figure 8.6: Electronegativities of Elements Electronegativity

12. Bond Polarity and Electronegativity Figure 8.7: Electronegativity and Bond Polarity There is no sharp distinction between bonding types. The positive end (or pole) in a polar bond is represented  + and the negative pole  -. HyperChem

13. Drawing Lewis Structures Follow Step by Step Method (See Ng Web-site)See Ng Web-site 1.Total all valence electrons. [Consider Charge] 2.Write symbols for the atoms and guess skeleton structure [ define a central atom ]. 3.Place a pair of electrons in each bond. 4.Complete octets of surrounding atoms. [ H = 2 only ] 5.Place leftover electrons in pairs on the central atom. 6.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. HyperChem CyberChem video

14. Lewis Structures - Examples

15. Exceptions to the Octet Rule 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 BF 3. Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

16. Exceptions – Central Atoms - Less than an Octet

17. Exceptions to the Octet Rule More than an Octet This is the largest class of exceptions. Atoms from the 3 rd period onwards can accommodate more than an octet. Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. HyperChem

18. Exceptions – Central Atoms - Greater than an Octet

19. There are five fundamental geometries for molecular shape: Molecular Shapes

20. Figure 9.3 HyperChem

21. Molecular Shapes – 3D Notations

22. Summary of VSEPR Molecular Shapes e-pairsNotationName of VSEPR shapeExamples 2AX 2 LinearHgCl 2, ZnI 2, CS 2, CO 2 3AX 3 Trigonal planarBF 3, GaI 3 AX 2 ENon-linear (Bent)SO 2, SnCl 2 4AX 4 TetrahedralCCl 4, CH 4, BF 4 - AX 3 E(Trigonal) PyramidalNH 3, OH 3 - AX 2 E 2 Non-Linear (Bent)H 2 O, SeCl 2 5AX 5 Trigonal bipyramidalPCl 5, PF 5 AX 4 EDistorted tetrahedral (see-sawed) TeCl 4, SF 4 AX 3 E 2 T-ShapedClF 3, BrF 3 AX 2 E 3 LinearI 3 -, ICl 2 - 6AX 6 OctahedralSF 6, PF 6 - AX 5 ESquare PyramidalIF 5, BrF 5 AX 4 E 2 Square PlanarICl 4 -, BrF 4 - See Ng Web-site HyperChemCyberChem video

23. Examples: VSEPR Molecular Shapes

24. The Effect of Nonbonding Electrons By experiment, the H-X-H bond angle decreases on moving from C to N to O: Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. Therefore, the bond angle decreases as the number of lone pairs increase VSEPR Model HyperChem

25. Figure 9.10: Shapes of Larger Molecules In acetic acid, CH 3 COOH, there are three central atoms. VSEPR Model HyperChem

26. Figure 8.10: Drawing Lewis Structures Resonance Structures

27. Figure 9.12 HyperChem

28. Figure 9.11: Molecular Shape and Molecular Polarity HyperChem

29. Figure 9.13: Molecular Shape and Molecular Polarity HyperChem

30. Lewis structures and VSEPR do not explain why a bond forms. How do we account for shape in terms of quantum mechanics? What are the orbitals that are involved in bonding? We use Valence Bond Theory: Bonds form when orbitals on atoms overlap. There are two electrons of opposite spin in the orbital overlap. Covalent Bonding and Orbital Overlap

31. Figure 9.14: Covalent Bonding and Orbital Overlap

34. To determine the electron pair geometry: draw the Lewis structure, count the total number of electron pairs around the central atom, arrange the electron pairs in one of the above geometries to minimize e  -e  repulsion, and count multiple bonds as one bonding pair. VSEPR Model (Figure 9.6)

35. VSEPR Model

38. Drawing Lewis Structures Formal Charge Consider: For C: There are 4 valence electrons (from periodic table). In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge: 4 - 5 = -1.

39. Drawing Lewis Structures Formal Charge Consider: For N: There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge = 5 - 5 = 0. We write: CyberChem video

40. Chemical Bonding Lewis AXE notation VSEPR shapes Polarity