Chemical Bonding Theory

1. CHEMICAL BONDING
•IONIC BONDS
•COVALENT BONDS
•HYDROGEN BONDS
•METALLIC BONDS

3. IONIC BONDING
When an atom of a nonmetal
takes one or more electrons
from an atom of a metal
so both atoms end up with
eight valence electrons

4. IONIC BONDING
IS THE COMPOUND
AN IONIC COMPOUND?
METAL
METAL NONMETAL
SUBSCRIPTS

5. IONIC BOND FORMATION
Neutral atoms come near each other.
Electron(s) are transferred from the Metal atom
to the Non-metal atom. They stick together
because of electrostatic forces, like magnets.

6. IONIC BONDING
Metals will tend to lose electrons
and become
POSITIVE CATIONS
Normal sodium atom loses one electron to become sodium ion

7. IONIC BONDING
Nonmetals will tend to gain
electrons and become
NEGATIVE ANIONS
Normal chlorine atom gains an electron to become a chloride ion

8. POLYATOMIC IONS--a group
of atoms that act like one ion
NH4
+1
--ammonium ion
CO3
-2
--carbonate ion
PO4
-3
--phosphate ion
IONIC BONDING

9. IONIC BONDING
SODIUM SULFATE

10. Properties of Ionic Compounds
• Crystalline structure.
• A regular repeating
arrangement of ions in the solid.
• Ions are strongly bonded.
• Structure is rigid.
• High melting points- because of
strong forces between ions.

11. Crystalline structure
+
+
+ +
+
+
+
+
+
-
-
-
-
- -
-
-
-
The
POSITIVE
CATIONS
stick to the
NEGATIVE
ANIONS, like
a magnet.

12. Do they Conduct?
• Conducting electricity is allowing
charges to move.
• In a solid, the ions are locked in place.
• Ionic solids are insulators.
• When melted, the ions can move
around.
• Melted ionic compounds conduct.
• First get them to 800ºC.
• Dissolved in water they conduct.

13. Ionic solids are brittle
+ - + -
+
- +
-
+ - + -
+
- +
-

14. Ionic solids are brittle
+ - + -
+
- +
-
+ - + -
+
- +
-
• Strong Repulsion breaks crystal apart.

16. COVALENT BONDING
When an atom of one
nonmetal
shares one or more
electrons
with an atom of another
nonmetal so both atoms
end up with
eight valence electrons

17. COVALENT BOND
FORMATION
When one nonmetal shares one or
more electrons with an atom of
another nonmetal so both atoms end
up with eight valence electrons

18. COVALENT BONDING
IS THE COMPOUND
A COVALENT COMPOUND?
NONMETAL
NONMETAL NONMETAL
YES since it is made of only nonmetal elements
YES since it is made of only nonmetal elements

19. Covalent bonding
• Fluorine has seven valence electrons
F

20. Covalent bonding
• Fluorine has seven valence electrons
• A second atom also has seven
F F

21. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
F F

22. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
F F

23. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
F F

24. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
F F

25. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
F F

26. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals
F F

27. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals
F F
8 Valence
electrons

28. Covalent bonding
 Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals
F F
8 Valence
electrons

29. Single Covalent Bond
• A sharing of two valence electrons.
• Only nonmetals and Hydrogen.
• Different from an ionic bond because they
actually form molecules.
• Two specific atoms are joined.
• In an ionic solid you can’t tell which atom
the electrons moved from or to.

30. Water
H
O
Each hydrogen has 1 valence
electron
Each hydrogen wants 1 more
The oxygen has 6 valence
electrons
The oxygen wants 2 more
They share to make each other
happy

31. Water
• Put the pieces together
• The first hydrogen is happy
• The oxygen still wants one more
H O

32. Water
• The second hydrogen attaches
• Every atom has full energy levels
H O
H

33. Carbon dioxide
• CO2 - Carbon is central atom (
I have to tell you)
• Carbon has 4 valence
electrons
• Wants 4 more
• Oxygen has 6 valence
electrons
• Wants 2 more
O
C

34. Carbon dioxide
• Attaching 1 oxygen leaves the oxygen 1
short and the carbon 3 short
O
C

35. Carbon dioxide
 Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2
short
O
C
O

36. Carbon dioxide
 The only solution is to share more
O
C
O

37. Carbon dioxide
 The only solution is to share more
O
C
O

38. Carbon dioxide
 The only solution is to share more
O
C
O

39. Carbon dioxide
 The only solution is to share more
O
C
O

40. Carbon dioxide
 The only solution is to share more
O
C
O

41. Carbon dioxide
 The only solution is to share more
O
C
O

42. Carbon dioxide
 The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
O
C
O

43. Carbon dioxide
 The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
O
C
O
8 valence
electrons

44. Carbon dioxide
 The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
O
C
O
8 valence
electrons

45. Carbon dioxide
 The only solution is to share more
 Requires two double bonds
 Each atom gets to count all the atoms in
the bond
O
C
O
8 valence
electrons

46. How to draw them
• Add up all the valence electrons.
• Count up the total number of electrons to
make all atoms happy.
• Subtract.
• Divide by 2
• Tells you how many bonds - draw them.
• Fill in the rest of the valence electrons to fill
atoms up.

47. Examples
• HCN C is central atom
• N - has 5 valence electrons wants 8
• C - has 4 valence electrons wants 8
• H - has 1 valence electrons wants 2
• HCN has 5+4+1 = 10
• HCN wants 8+8+2 = 18
• (18-10)/2= 4 bonds
• 3 atoms with 4 bonds -will require multiple bonds -
not to H

48. HCN
• Put in single bonds
• Need 2 more bonds
• Must go between C and N
N
H C

49. HCN
 Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
N
H C

50. HCN
 Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
 Must go on N to fill octet
N
H C

51. Polar Bonds
• When the atoms in a bond are the same, the
electrons are shared equally.
• This is a nonpolar covalent bond.
• When two different atoms are connected,
the atoms may not be shared equally.
• This is a polar covalent bond.
• How do we measure how strong the atoms
pull on electrons?

52. Electronegativity
• A measure of how strongly the atoms attract
electrons in a bond.
• The bigger the electronegativity difference
the more polar the bond.
• 0.0 - 0.3 Covalent nonpolar
• 0.3 - 1.67 Covalent polar
• >1.67 Ionic

53. How to show a bond is polar
• Isn’t a whole charge just a partial charge
means a partially positive
means a partially negative
• The Cl pulls harder on the electrons
• The electrons spend more time near the Cl
H Cl
 

54. Polar Molecules
Molecules with ends

55. Polar Molecules
• Molecules with a positive and a negative end
• Requires two things to be true
 The molecule must contain polar bonds
This can be determined from differences in
electronegativity.
Symmetry can not cancel out the effects of the
polar bonds.
 Must determine geometry first.

56. Is it polar?
• HF
• H2O
• NH3
• CCl4
• CO2

57. Intermolecular Forces
What holds molecules to each other

58. Intermolecular Forces
• They are what make solid and liquid molecular
compounds possible.
• The weakest are called van der Waal’s forces -
there are two kinds
• Dispersion forces
• Dipole Interactions
– depend on the number of electrons
– more electrons stronger forces
– Bigger molecules

59. Dipole interactions
• Depend on the number of electrons
• More electrons stronger forces
• Bigger molecules more electrons
•Fluorine is a gas
•Bromine is a liquid
•Iodine is a solid

60. Dipole interactions
• Occur when polar molecules are attracted to
each other.
• Slightly stronger than dispersion forces.
• Opposites attract but not completely hooked
like in ionic solids.

61. Dipole interactions
• Occur when polar molecules are attracted to
each other.
• Slightly stronger than dispersion forces.
• Opposites attract but not completely hooked
like in ionic solids.
H F


H F



62. Dipole Interactions

























63. Hydrogen bonding
• Are the attractive force caused by hydrogen
bonded to F, O, or N.
• F, O, and N are very electronegative so it is
a very strong dipole.
• The hydrogen partially share with the lone
pair in the molecule next to it.
• The strongest of the intermolecular forces.

64. Hydrogen Bonding
H
H
O
+
-
+
H H
O
+
-
+

65. Hydrogen bonding
H
H
O H
H
O
H
H
O
H
H
O
H
H
O
H
H
O
H
H
O

66. MOLECULAR
SHAPES
OF
COVALENT
COMPOUNDS

67. VSepR tHEORY

68. 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.

69. 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.

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

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

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

74. Bent #2
•Number of Bonds = 2
•Number of Shared Pairs of Electrons = 2
•Number of Unshared Pairs of Electrons = 1
•Bond Angle = >120°
EXAMPLE:
O3

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

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

77. Trigonal bIPyramidal
•Number of Bonds = 5
•Number of Shared Pairs of Electrons = 5
•Number of Unshared Pairs of Electrons = 0
•Bond Angle = <120°
EXAMPLE:
NbF5

78. OCTAHEDRAL
•Number of Bonds = 6
•Number of Shared Pairs of Electrons = 6
•Number of Unshared Pairs of Electrons = 1
•Bond Angle = 90°
EXAMPLE:
SF6

79. Metallic Bonds
• How atoms are held together in
the solid.
• Metals hold onto there valence
electrons very weakly.
• Think of them as positive ions
floating in a sea of electrons.

80. Sea of Electrons
+ + + +
+ + + +
+ + + +
• Electrons are free to move through the
solid.
• Metals conduct electricity.

81. Metals are Malleable
• Hammered into shape
(bend).
• Ductile - drawn into
wires.

82. Malleable
+ + + +
+ + + +
+ + + +

83. Malleable
+ + + +
+ + + +
+ + + +
• Electrons allow atoms to slide by.