1. Bonding
Chapter 8

2. Two general types of bonds
Ionic: the transfer of electrons from a metal to a non metal.
Covalent: the sharing of electrons between atoms.
Valence electrons: the electrons being transferred or shared

3. Valence Electrons
The electrons used to form bonds

4. Valence electrons are depicted using Lewis Dot Structures

5. Ionic Bonding

6. Covalent Bonding A type of bonding between non-metals.
This involves the sharing of electrons by 2 or more atoms.
The electrons are shared, not transferred as with ionic bonds.

8. When nonmetals bond together a covalent bond is created and we call them molecules or molecular compounds!

9. Molecules
Molecules are neutral atoms that are joined together by covalent bonds
Molecular formula - shows you how many atoms of each element is in a substance
Example: CO2 , NH4

10. Octet Rule and Covalent Bonding
An octet is 8 valence electrons around an atom to achieve a noble gas configuration!
Molecules want the same thing, but they share their valence electrons to achieve the octet rule.

11. EXAMPLES INCLUDE THE BONDS BETWEEN:
H F2 Br2 Cl2 HCl H2O
Other than hydrogen these elements are on the right side of the periodic table.

12. Cl Cl Lewis Dot Structures
Formulas in which atomic symbols represent the element and all inner-shell electrons, dots represent valence electrons and dashes between two atomic symbols represent electron pairs in covalent bonds. Cl Cl

13. Single Covalent Bonds
When atoms share one pair of electrons they form a single covalent bond

14. Chlorine
forms a
covalent
bond
with
itself
Cl2

15. How
will
two
chlorine
atoms
react? Cl Cl

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

17. Cl Cl
what can they do to achieve an octet?
Share unpaired electrons!

18. Cl Cl

19. Cl Cl

20. Cl Cl

21. Cl Cl

22. Cl Cl
octet

23. Cl Cl
octet

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

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

26. Cl Cl
This is the bonding pair
(shared pair of electrons)

27. Cl Cl
It is a single bonding pair

28. Cl Cl
It is called a SINGLE BOND

29. Single bonds are abbreviatedCl Cl
Single bonds are abbreviated
with a dash
Normally in the final structure the
valence electrons are not drawn

30. This is the chlorine molecule,Cl Cl
This is the chlorine molecule,
Cl2

31. Neutral Molecules
4 bonds
3 bonds
2 bonds
1 bond

32. Double Bonds
Sharing of two pairs of
electrons between two atoms

33. O2
Oxygen is also one of the diatomic molecules

34. O
How will two oxygen atoms bond?

35. O
Each atom has two unpaired electrons

36. O

37. O

38. O

39. O

40. O

41. O

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

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

44. O

45. O O

46. O O

47. O O

48. Both electron pairs are shared.

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

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

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

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

53. This is the oxygen molecule,=
This is the oxygen molecule, O2

54. Triple Bonds
Atoms that share three pairs of electrons:
Example: N2

55. Molecular Shapes

59. bond polarity Covalent bonds can be polar or non-polar.
If a bond is non-polar, that means that there is an equal sharing of electrons between atoms (Cl2).
If a bond is polar, that means the electrons are not shared equally, making one side of the bond more negative and the other side more positive.
The electrons are more strongly attracted to the more electronegative atom.

61. The dipole’s direction is from the positive pole to the negative pole.
Dipole – is created by opposite charges that are separated by a short distance.
The dipole’s direction is from the positive pole to the negative pole.
The dipole (magnet) is created due to electronegativity differences between atoms.
H Cl
(H= 2.1; Cl = 3.0)

64. Electronegativity 0.0-0.3 Non polar covalent 0.3-1.7 polar covalent
>/= 1.8 Ionic

65. Polar Compounds Need to look at all the bonds in the compound.
Are individual bonds polar or non polar.
What is the sum of all the bonds.
atom. electronegativity

70. Bond Lengths
and
Bond Strengths

71. Multiple Bonds
Double bonds are stronger and shorter than single bonds.
Triple bonds are even stronger and shorter than both double and single bonds.
Common between carbon, oxygen and nitrogen atoms.

72. Bond Lengths and Bond Energies