Chapter 6

1. Chapter 6 Covalent Compounds

2. 6.1 Covalent Bonds • Sharing Electrons • Covalent bonds form when atoms share one or more pairs of electrons • nucleus of each atom is attracted to electron cloud of other atom • neither atom removes an electron from the other

3. Covalent Bonding

4. 6.1 Covalent Bonds • Sharing Electrons • Covalent bonds • space where electrons move is called molecular orbital • made when atomic orbitals overlap

5. Molecules

6. 6.1 Covalent Bonds • Energy and Stability • Noble gases are stable (full octet) (low P.E.) • Other elements are not stable (high P.E.) • covalent bonding decreases potential energy because each atom achieves electron configuration like noble gas

7. 6.1 Covalent Bonds • Energy and Stability • because P.E. decreases when atoms bond, energy is released • i.e., atoms lose P.E. when they bond • loss of P.E. implies higher stability

9. 6.1 Covalent Bonds • Energy and Stability • potential energy determines bond length • at minimum P.E., distance between two bonded atoms is called bond length • bonded atoms vibrate • therefore, bond length is an average length

10. 6.1 Covalent Bonds • Energy and Stability • bonds vary in strength • bond energy is the amount of energy required to break the bonds in 1 mol of a chemical compound • bond energy predicts reactivity • bond energy is equal to loss of P.E. during formation

11. Bond Energies and Lengths

12. 6.1 Covalent Bonds • Electronegativity • Atoms share electrons equally or unequally • nonpolar covalent bond: bonding electrons shared equally • polar covalent bond: shared electrons more likely to be found around more electronegative atom

13. 6.1 Covalent Bonds • Electronegativity • Atoms share electrons equally or unequally • difference in electronegativity can be used to predict type of bond (but boundaries are arbitrary) • I think this concept is important for AP Biology.

15. Bond Types

16. 6.1 Covalent Bonds • Electronegativity • Polar molecules have positive and negative ends • such molecules called dipoles •  (“delta”) means partial in math and science • positive end—+ • negative end—- • example: H+F-

17. Electronegativity Difference for Hydrogen Halides

18. 6.1 Covalent Bonds • Electronegativity • Polarity is related to bond strength • greater electronegativity means • greater polarity means • greater bond strength

19. 6.1 Covalent Bonds • Electronegativity • Bond type determines properties of substances • metallic bonds: electrons can move from one atom to another—good conductors • ionic bonds: hard and difficult to break apart • covalent bonds: low melting/boiling points

20. Properties of Substances with Different Types of Bonds

21. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Lewis structures represent valence electrons with dots • position of electrons is symbolic (not literal) • shows only the valence electrons of an atom • dots around atomic symbol represent electrons

22. Lewis Structures of Second-Period Elements

23. Lewis Structures of Second-Period Elements

24. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Cl2 • HCl

25. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Drawing • 1. Gather information • draw Lewis structure for each atom in compound; place one electron on each side before pairing • determine total number of valence electrons

26. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Drawing • 2. Arrange atoms • arrange structure to show bonding • halogens and hydrogen usually make one bond at end of molecule • carbon usually in center

27. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Drawing • 3. Distribute the dots so that each atom satisfies octet rule (except H, Be, B) • 4. Draw the bonds as long dashes • 5. Verify the structure by counting number of valence electrons

28. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Polyatomic Ions • use brackets [] to show overall charge • example:

29. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Multiple Bonds • sharing two pairs of electrons is a double bond • sharing three pairs of electrons makes triple bonds • example:

30. 6.2 Drawing and Naming • Lewis Electron-Dot Structures • Resonance Structures • sometimes, multiple structures are possible • show all possibilities • example:

31. 6.2 Drawing and Naming • Naming Covalent Compounds • First name: name of first element in formula • usually least electronegative • requires a prefix if more than one of them • Second name: ends in –ide • requires a prefix if more than one of them

32. Naming Covalent Compounds

33. 6.3 Molecular Shapes • Determining Molecular Shapes • Three-dimensional shape helps determine physical and chemical properties • valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes • based on idea that electrons repel one another

34. Molecular Shapes

35. 6.3 Molecular Shapes • Determining Molecular Shapes • Let’s try some. • CO • CO2 • BF3 • CH4 • SnCl2 • SO2

36. Simple Shapes

37. Trigonal Planar

38. Tetrahedral

39. Bent

40. 6.3 Molecular Shapes • Molecular Shape Affects Properties • Shape affects polarity • compare CO2 and H2O • polarity affects properties (such as boiling point) due to attractions between molecules

41. Polar Bonds