BONDING

1. BONDING PART II

2. LET’S FIRST REVIEW IONIC BONDING

3. K F In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds.

4. _ + K F The compound potassium fluoride consists of potassium (K+) ions and fluoride (F-) ions

5. _ + K F The ionic bond is the attraction between the positive K+ ion and the negative F- ion An ionic bond is a metal and nonmetal

6. So what are covalent bonds?

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

8. 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.

9. 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

10. 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 A covalent bond is between a nonmetal and a nonmetal

11. Chlorine forms a covalent bond with itself Cl2

12. How will two chlorine atoms react? Cl Cl

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

14. 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?

15. Cl Cl The octet is achieved by each atom sharing the electron pair in the middle

16. Cl Cl The octet is achieved by each atom sharing the electron pair in the middle

17. Cl Cl This is the bonding pair

18. Cl Cl It is a single bonding pair

19. Cl Cl It is called aSINGLE BOND

20. Cl Cl Single bonds are abbreviated with a dash

21. Cl Cl This is the chlorine molecule, Cl2

22. Diatomicatoms • These are atoms that are never, never, never found alone in nature. • H2, N2, O2, F2, Cl2, Br2, I2

23. Remember the number 7 • There are 7 diatomic molecules

24. Start with 7 on the periodic table • That would be Nitrogen (this is one of the diatomic molecules)

25. Then make a number 7 on the periodic table

26. Do not forget the 7th diatomic molecule

27. Naming covalent bonds

28. Rules: 1. The formula is written with the more electropositive element (the one further to the left on the periodic table) placed first, then the more electronegative element (the one further to the right on the periodic table). Example : CO2 [Important exception: when the compound contains oxygen and a halogen, the halogen is placed first. If both elements are in the same group, the one with the higher period number is named first.] Example: Cl2O .

29. 2. The second element’s name ends in –ide 3. Never use the prefix mono with the 1st word. 4. Do not put two O’s together

30. Naming covalent bonds • Prefixes are used for the number of atoms • 1 – mono • 2 – di • 3 – tri • 4 – tetra • 5 – penta • 6 – hexa • - hepta • - octa • - nona • - deca

31. Examples: CO2 – CO – P4Cl3 - Carbon Dioxide Carbon monoxide Tetraphosphorous trichloride

32. O O How will two oxygen atoms bond?

33. O2 Oxygen is one of the diatomic molecules

34. O O Each atom has two unpaired electrons

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

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

37. O O two bonding pairs, making a doublebond

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

39. Triple bonds are when three sets of electrons are shared.

40. Some covalent bonds can have more than one bond type.

41. MOLECULAR SHAPES OF COVALENT COMPOUNDS

42. VSepR tHEORY ALENCE HELL VSEPR LECTRON AIR EPULSION

43. What Vsepr means Since electrons do not like each other, because of their negative charges, they orient themselves as far apart as possible, from each other. This leads to molecules having specific shapes.

44. Things to remember • Atoms bond to form an Octet (8 outer electrons/full outer energy level) • Bonded electrons take up less space then un-bonded/unshared pairs of electrons.

45. HERE ARE THE RESULTING MOLECULAR SHAPES

46. Linear EXAMPLE: BeF2 • Number of Bonds = 2 • Number of Shared Pairs of Electrons = 2 • Bond Angle = 180°

47. Bent #1 EXAMPLE: H2O • Number of Bonds = 2 • Number of Shared Pairs of Electrons = 2 • Number of Unshared Pairs of Electrons = 2 • Bond Angle = < 120°

48. Trigonal Planar EXAMPLE: GaF3 • Number of Bonds = 3 • Number of Shared Pairs of Electrons = 3 • Number of Unshared Pairs of Electrons = 0 • Bond Angle = 120°

49. Trigonal Pyramidal EXAMPLE: NH3 • Number of Bonds = 3 • Number of Shared Pairs of Electrons = 4 • Number of Unshared Pairs of Electrons = 1 • Bond Angle = <109.5°

50. Tetrahedral EXAMPLE: CH4 • Number of Bonds = 4 • Number of Shared Pairs of Electrons = 4 • Number of Unshared Pairs of Electrons = 0 • Bond Angle = 109.5°