Chemical Bonding

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 • ClO2 • FeSO4 • H2O • NaNO3

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 • CaCl2 • Be3N2

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 Na Ca S Cl

16. Lewis Formulas for Diatomics

17. Valence Electrons

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

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 • H3CCN • NH2OH

22. Resonance

23. PROBLEM • Nitric Oxide, N2O, NNO arrangement.

24. Exceptions to the Octet Rule • Odd Number of Electrons (NO) • Unfilled Octet (BH3) • 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 (AX2, AX3, AX4)

29. EXAMPLE • CH4 • NH3 • H20

30. PROBLEM • Determine the shape. • CO32- • SCl2

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.