1. Covalent Bonding Chapter 8

3. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Atom Oxygen Molecule (O 2 )

8. Covalent Bonds

9. Chapter 2 Chemical Principles Bonding Covalent bonding

10. So what are covalent bonds?

11. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule).

12. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair.

13. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair. The shared electron pair is called a bonding pair

14. Cl 2 Chlorine forms a covalent bond with itself

15. Cl Cl How will two chlorine atoms react?

16. Cl Cl Each chlorine atom wants to gain one electron to achieve an octet

17. Cl Cl Neither atom will give up an electron – chlorine is highly electronegative. What’s the solution – what can they do to achieve an octet?

18. Cl Cl Neither atom will give up an electron – chlorine is highly electronegative. What’s the solution – what can they do to achieve an octet?

19. Cl Cl octet

20. Cl Cl circle the electrons for each atom that completes their octets octet

21. Cl Cl circle the electrons for each atom that completes their octets The octet is achieved by each atom sharing the electron pair in the middle

22. Cl Cl circle the electrons for each atom that completes their octets The octet is achieved by each atom sharing the electron pair in the middle

23. Cl Cl circle the electrons for each atom that completes their octets This is the bonding pair

24. Cl Cl circle the electrons for each atom that completes their octets It is a single bonding pair

25. Cl Cl circle the electrons for each atom that completes their octets It is called a SINGLE BOND

26. Cl Cl circle the electrons for each atom that completes their octets Single bonds are abbreviated with a dash

27. Cl Cl circle the electrons for each atom that completes their octets This is the chlorine molecule, Cl 2

28. O 2 Oxygen is also one of the diatomic molecules

29. How will two oxygen atoms bond? O O

30. Each atom has two unpaired electrons O O

31. O O

32. O O

33. Oxygen atoms are highly electronegative. So both atoms want to gain two electrons. O O

34. O O Both electron pairs are shared.

35. 6 valence electrons plus 2 shared electrons = full octet O O

36. 6 valence electrons plus 2 shared electrons = full octet O O

37. two bonding pairs, O O making a double bond

38. For convenience, the double bond can be shown as two dashes. O O = O O

39. This is the oxygen molecule, O 2 O O = this is so cool!!

40. Multiple Covalent bonds O O Sharing One Pair of electrons One Covalent Bond Only 7 electrons does Not meet Octet Rule! Need to share Another pair of electrons O O Sharing Two Pairs of electrons Two Covalent Bonds A Double Bond A Double Bond can be represented by a double line O O

41. Multiple Covalent bonds N N N N Sharing Three Pairs of electrons Three Covalent Bonds A Triple Bond A Triple Bond can be represented by a Triple line Nitrogen

45. Drawing Lewis Dot Structures Draw Lewis Dot Structures for: PH 3 H 2 S HCl CCl 4 SiH 4 CH 2 Cl 2

52. INTRODUCTION A) Lewis structures do not indicate the three dimensional shape of a molecule. They do not show the arrangement space of the atoms, what we call the molecular geometry or molecular structure. B) Molecules have definite shapes and the shape of a molecule controls some of its chemical and physical properties.

53. II. Valence Shell Electron Pair Repulsion Theory - VSEPR - predicts the shapes of a number of molecules and polyatomic ions. A) Assumptions of VSEPR Theory 1) Electron pairs in the valence shell of an atom tend to orient themselves so that the total energy is minimized. This means that: the electrons will approach the nucleus as close as possible yet take positions as far away from each other as possible to minimize _______________ .

54. 2) Because lone pairs of electrons are spread out more broadly than bond pairs, repulsions are greatest between two lone pairs, intermediate between a lone pair and a bond pair , and weakest between two bonding pairs of electrons. 3) Repulsive forces decrease rapidly with increasing interpair angle - greatest at 90 o , much weaker at 120 o , and very weak at 180 o . B) What are the ideal arrangements of electron pairs to minimize repulsions?

56. Double and even triple bonds are commonly observed for C, N, P, O, and S H 2 CO SO 3 C 2 F 4

57. Some Common Geometries Linear Trigonal Planar Tetrahedral

67. Polar Covalent Bonds: Unevenly matched, but willing to share.

71. Dipole-Dipole Two polar molecules align so that  + and  - are matched (electrostatic attraction) Ex: ethane (C 2 H 6 ) vs. fluromethane (CH 3 F) Occurs when polar molecules are attracted to one another. The slightly region of a polar molecule is weakly attracted to the slightly positive region of another polar molecule. Similar to but much weaker than ionic bonds.