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

24. 6.2 Drawing and Naming  Lewis Electron-Dot Structures  Cl 2  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  CO 2  BF 3  CH 4  SnCl 2  SO 2

36. Simple Shapes

37. Trigonal Planar

38. Tetrahedral

39. Bent

40. 6.3 Molecular Shapes  Molecular Shape Affects Properties  Shape affects polarity  compare CO 2 and H 2 O  polarity affects properties (such as boiling point) due to attractions between molecules

41. Polar Bonds