1. Chemical Bonding Chapter 8

2. 8.1 Types of Bonds

3. Ionic and Covalent Bonds Chemical Bonds are the force that holds atoms together in a compound or molecule Ionic Bonds are bonds between ions held together by electrostatic force. Covalent Bonds are where electrons are shared between atoms.

4. PROBLEM Identify which type of Chemical Bond is most likely in the following. NaF ClO 2 FeSO 4 H 2 O NaNO 3

5. Polar and Nonpolar Covalent Bonds Rarely are covalent bonds a completely equal sharing. Polarity is the degree of transfer from one member of a covalent bond to the other. Nonpolar Covalent Bonds have little to no polarity and share the electrons equally. Polar Covalent Bonds have detectable polarity and have an uneven sharing of electrons.

6. Polarity Continuum

7. Electronegativity Electronegativity – The ability for an atom to bonding electrons. Varies in a periodic fashion The greater the difference in electronegativity the greater the polarity of the bond.

8. Electronegativity

9. 8.2 Ionic Bonds

10. Lewis Symbols Lewis Symbols – A shorthand to show the valence electrons of an atom. Examples

11. PROBLEM Draw the Lewis Structures of the following. NaF CaCl 2 Be 3 N 2

12. Structure of Ionic Crystals An ionic crystal is an arrangement of ions that maximize attraction and minimize repulsion of ions. The crystals structure makes ionic solids very hard, brittle and poor conductors. Crystal structure also accounts for high melting and boiling temperatures.

13. 8.3 Covalent Bonds

14. The Octet Rule In the Main Group elements, stability is reached by becoming isoelectronic with the noble gases. This completes the Valance Shell for the principal energy level Since the s and p orbitals take eight electrons, this is called the octet rule.

15. Octet Examples NaCaSCl

16. Lewis Formulas for Diatomics

17. Valence Electrons

18. PROBLEM Give Possible Identities for each X. XXX Cl

19. Structures of Covalent Molecules Write the Skeleton Equation Sum the Valance Electrons and determine the total Place two electrons for each single bond If you have remaining valance electrons, add them as unshared pairs to satisfy unfilled octets Use double and triple bonds to satisfy octets on the central atoms

20. EXAMPLE Carbon Dioxide

21. PROBLEM H 3 CCN NH 2 OH

22. Resonance

23. PROBLEM Nitric Oxide, N 2 O, NNO arrangement.

24. Exceptions to the Octet Rule Odd Number of Electrons (NO) Unfilled Octet (BH 3 ) More than eight electrons around the central atom.

25. Bonding in Carbon Compounds Carbon’s versatility comes from its four valance electrons. Carbon can readily bond with itself at “normal” temperatures

26. 8.4 Shape of Molecules

27. The VSEPR Theory Valence Shell Electron Pair Repulsion Theory says that pairs of electrons will try to get as far away from each other as possible. You use the Lewis structure to determine a general structure then fine tune the model.

28. Predicting the Shape of Molecules Draw the Lewis formula Count the number of bonds and unshared pairs on the central atom The sum gives you the parent formula (Linear, Trigonal Planar, Tetrahedral) Consider on the bonded atoms to determine the sub-shape (AX 2, AX 3, AX 4 )

29. EXAMPLE CH 4 NH 3 H 2 0

30. PROBLEM Determine the shape. CO 3 2- SCl 2

31. Polarity of Molecules If the bonds are polar and the molecular shape is not symmetrical the molecule is polar. If the bonds are not polar or the molecule is symmetrical the molecule is nonpolar.