Chapter 8

1. Chapter 8 Covalent Bonding

2. Section 1 Molecular Compounds

3. Section 1 Learning Targets 8.1.1 – I can distinguish between the melting points and boiling points of molecular compounds and ionic compounds. 8.1.2 – I can describe the information provided by a molecular formula.

4. Molecules and Molecular Compounds • Covalent bond – atoms held together by sharing electrons. • Molecule – neutral group of atoms joined together by covalent bonds.

5. Molecular compound – compound composed of molecules. • Diatomic molecule – molecule consisting of two atoms.

6. Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

7. Molecular Formulas • Molecular formula – chemical formula of a molecular compound. • A molecular formula shows how many atoms of each element a molecule contains.

9. Molecular formulas can not tell you about a molecules structure.

10. Section 2 The Nature of Covalent Bonding Part 1

11. Section 2 Learning Targets 8.2.1 – I can describe how electrons are shared to form covalent bonds and identify exceptions to the octet rule. 8.2.2 – I can demonstrate how electron dot structures represent shared electrons. 8.2.3 – I can describe how atoms form double or triple bonds.

12. Section 2 Learning Targets 8.2.4 – I can distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy. 8.2.5 – I can describe how oxygen atoms are bonded in ozone.

13. The Octet Rule in Covalent Bonding • In covalent bonding, electron sharing usually occurs so that atoms can attain the electron configuration of the noble gases.

14. Single Covalent Bonds • Structural formula – represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms.

15. An electron dot structure, such as H:H represents the shared pair of electrons of the covalent bond by two dots.

16. Single covalent bond – two atoms held together by sharing a pair of electrons. • Unshared pair – (lone pair) or nonbonding pair of electrons.

17. In the water molecule the two hydrogen atoms share electrons with the one oxygen to attain a noble gas electron configuration.

18. The halogens form single covalent bonds because they have seven valence electrons and need one more to attain the electron configuration of a noble gas.

19. Ammonia has three bonds and one unshared pair of electrons.

20. Carbon behaves differently than expected. • What you expect is: • What happens is: • This movement of electrons allows carbon to make four covalent bonds.

22. Double and Triple Covalent Bonds • Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two or three pairs of electrons. • Double covalent bond – bond that involves two shared pairs of electrons. • Triple covalent bond – bond formed by sharing three pairs of electrons.

23. Oxygen forms a double bond by sharing two pairs of electrons.

24. Nitrogen’s three 2p electrons allow it to form triple bonds.

25. Some elements exist as diatomic molecules. • There are seven of them.

26. Double and triple bonds can exist in molecules that are not diatomic.

27. Section 2 The Nature of Covalent Bonding Part 2

28. Coordinate Covalent Bonds • Coordinate covalent bond – covalent bond in which one atom contributes both bonding electrons. • Represented by an arrow pointing to the atom that accepts the electrons.

29. In the CO molecule the oxygen is stable with the double bond but the carbon is not.

30. The problem is solved when oxygen donates a pair of electrons to the bond.

31. Polyatomic ions – a tightly bound group of atoms that has a positive or negative charge and behaves like a unit.

32. Polyatomic ions usually contain both covalent and coordinate bonds.

34. Bond Dissociation Energy • Bond dissociation energy – energy required to break the bond between two covalently bonded atoms. • A large bond dissociation energy corresponds to a strong covalent bond.

35. Resonance • Resonance structures – structure that occurs when it is possible to draw two or more valid electron dot structures. • Usually seen with double bonds – where they could shift around.

36. The actual bonding in ozone (O3) is a hybrid, mixture, of the extremes represented by the resonance forms.

37. Exceptions to the Octet Rule • The octet rule can not be satisfied in molecules whose total valence electrons in an odd number. • There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons.

38. Some boron compounds have an unfilled octet.

39. Sulfur and phosphorus compounds expand to more than an octet. video

40. Section 3 Bonding Theories

41. Section 3 Learning Targets 8.3.2 – I can describe how VSEPR theory helps predict the shapes of molecules.

42. VSEPR Theory • VSEPR (Valence Shell Electron Pair Repulsion Theory • VSEPR video • According to VSEPR the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.

43. Unshared pairs are just as important as bonds because they help determine the shapes of the molecules. • Unshared pairs of electrons are held closer to the nucleus and push bonded atoms out of their way. • Tetrahedral angle – a bond angle of 109.5° that results when a central atom forms four bonds.

45. I’m giving you a handout instead of this table for your notes

48. Let’s Practice • On the back of your yellow paper answer the following:

49. Draw the molecules and tell their molecular shape • CO • BF3 Shape = Shape =

50. Draw the molecules and tell their molecular shape • SO2 • CH4 Shape = Shape =