1. Mr. Matthew Totaro Legacy High School Honors Chemistry
Chemical
Bonding
Mr. Matthew Totaro
Legacy High School
Honors Chemistry

2. "Perhaps one of you gentlemen would mind telling me just what it is outside the window that you find so attractive...?"

3. Bonding Theories Bonding is the way atoms attach to make molecules.
A chemical bond occurs when valence electrons are either transferred or shared between the nuclei of two atoms.

4. Sodium to Chlorine Bond
Chemical Bonds
Sodium to Chlorine Bond

5. Types of Bonds

6. Types of Bonds
We can classify bonds based on the kinds of atoms that are bonded together
Types of Atoms
Type of Bond
Bond Characteristic
metals to
nonmetals
Ionic
electrons
transferred
nonmetals to
Covalent
shared
metals
Metallic
pooled

7. Bond Polarity
The larger the difference in how strong one atom pulls on the valence electrons, the more polar the bond is.
Negative end toward more stronger atom.
d+ H — F d-

8. Bond Polarity
Bonding between unlike atoms results in unequal sharing of the electrons.
One atom pulls the electrons in the bond closer to its side.
One end of the bond has larger electron density than the other.
The result is bond polarity.
The end with the larger electron density gets a partial negative charge and the end that is electron deficient gets a partial positive charge.
H Cl • d+ d-

9. Types of Covalent Bonds
Nonpolar (Pure) Covalent Bond
e- are shared equally
symmetrical e- density
usually identical atoms

10. Types of Covalent Bonds
Polar Covalent Bond
e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

11. Bond Polarity
Nonpolar
Polar
Ionic .

12. Electronegativity
Measure of the pull an atom has on bonding electrons.
Increases across the period (left to right).
Decreases down the group (top to bottom).
The larger the difference in electronegativity, the more polar the bond.
Negative end toward more electronegative atom.

13. Electronegativity, Continued
2.1
1.0
0.9
0.8
0.7
1.5
1.2
1.3
1.1
1.4
1.6
1.8
1.7
1.9
2.2
2.4
2.0
2.5
3.0
3.5
4.0
2.8

14. Electronegativity, Continued

15. Electronegativity and Bond Polarity
If the difference in electronegativity between bonded atoms is 0 to 0.49, the bond is pure (nonpolar) covalent.
If the difference in electronegativity between bonded atoms 0.5 to 1.7, the bond is polar covalent.
If the difference in electronegativity between bonded atoms larger than 1.7, the bond is ionic. 15

16. Bond Polarity 3.0-3.0 = 0.0 3.0-2.1 = 0.9 3.0-1.0 = 2.0 Covalent Ionic
Nonpolar
Polar
0.49
1.7
4.0
Electronegativity difference

17. Dipole Moments
A dipole is a material with positively and negatively charged ends.
Polar bonds or molecules have one end slightly positive, d+, and the other slightly negative, d-.
Not “full” charges, come from nonsymmetrical electron distribution.
Dipole moment, m, is a measure of the size of the polarity.
Measured in debyes, D.

18. For Each of the Following Bonds, Determine Whether the Bond Is Ionic or Covalent. If Covalent, Determine if It Is Polar or Pure. If Polar, Indicate the Direction of the Dipole.
Pb-O
P-S
Mg-Cl
H-O

19. For Each of the Following Bonds, Determine Whether the Bond Is Ionic or Covalent. If Covalent, Determine if It Is Polar or Pure. If Polar, Indicate the Direction of the Dipole, Continued.
Pb-O ( ) = 1.6 \ polar covalent.
P-S ( ) = 0.4 \ nonpolar covalent.
Mg-Cl ( ) = 1.8 \ ionic.
H-O ( ) = 1.4 \ polar covalent.

20. Lewis Theory
Lewis bonding theory emphasizes the importance of valence electrons.
Uses dots to represent valence electrons either on or shared by atoms.
Arranges bonding between atoms to attain certain sets of stable valence electron arrangements.
G.N. Lewis
( )

21. Lewis Symbols of Atoms Also known as electron dot symbols.
Uses symbol of element to represent nucleus and inner electrons.
Uses dots around the symbol to represent valence electrons.
Put 2 electrons on one side, then one on the other 3.
Remember that elements in the same group have the same number of valence electrons; therefore, their Lewis dot symbols will look alike.

22. Lewis Symbols of Ions
Cations have Lewis symbols without valence electrons.
Lost in the cation formation.
They now have a full “outer” shell that was the previous second highest energy energy level.
Anions have Lewis symbols with 8 valence electrons.
Electrons gained in the formation of the anion.
Li• Li+ :F: [:F:]− • ••

23. Practice—Write the Lewis Symbol for Arsenic.

24. 9.1

25. Octet Rule -atoms will tend to gain, lose, or share
electrons until their outer energy level contains eight electrons (Noble Gas e- configuration), or two like Helium
-this produces the maximum stability in an atom

26.                
 + 
                           
     
sodium metal
chlorine gas
table salt

27. Lewis Bonding Theory
Atoms bond because it results in a more stable electron configuration.
Atoms bond together by either transferring or sharing electrons.
Usually this results in all atoms obtaining an outer energy level with 8 electrons.
Octet rule.
There are some exceptions to this rule—the key to remember is to try to get an electron configuration like a noble gas.
H, Li & Be try to achieve the He electron arrangement.

28. Ionic Bonds Metal to nonmetal. Metal loses electrons to form cation.
Nonmetal gains electrons to form anion.
Ionic bond results from + to − attraction.
Lewis theory allows us to predict the correct formulas of ionic compounds.

29. Predict the formula of the compound that forms between
Example—Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds
Predict the formula of the compound that forms between
calcium and chlorine. Cl ∙ Ca ∙
Draw the Lewis dot symbols
of the elements. Cl ∙ Cl ∙
Transfer all the valance electrons
from the metal to the nonmetal,
adding more of each atom as you
go, until all electrons are lost
from the metal atoms and all
nonmetal atoms have 8 electrons. Ca ∙
Ca2+
CaCl2

30. Practice—Use Lewis Symbols to Predict the Formula of an Ionic Compound Made from Reacting a Metal, Mg, that Has 2 Valence Electrons with a Nonmetal, N, that Has 5 Valence Electrons.

31. Practice—Use Lewis Symbols to Predict the Formula of an Ionic Compound Made from Reacting a Metal, M, that Has 2 Valence Electrons with a Nonmetal, X, that Has 5 Valence Electrons, Continued.
Mg3N2

32. Covalent Bonds Often found between two nonmetals.
Atoms bonded together to form molecules.
Strong attraction.
Atoms share pairs of electrons to attain octets (or duets).
Molecules generally weakly attracted to each other.
Observed physical properties of molecular substance due to these attractions.

33. Single Covalent Bonds •• • •• • • • • •• •• •• •• •• •• F H O H F H O
Two atoms share one pair of electrons.
2 electrons.
One atom may have more than one single bond. F •• • •• H • • O • • H •• F •• •• H •• O •• H •• F

34. Double Covalent Bond •• • •• • ••
Two atoms sharing two pairs of electrons.
4 electrons.
Shorter and stronger than single bond. O •• • O •• • O •• O

35. Triple Covalent Bond •• • •• • ••
Two atoms sharing 3 pairs of electrons.
6 electrons.
Shorter and stronger than single or double bond. N •• • N •• • N •• N

36. Bonding and Lone Pair Electrons
Electrons that are shared by atoms are called bonding pairs.
Electrons that are not shared by atoms but belong to a particular atom are called lone pairs.
Also known as nonbonding pairs.
O S O
Bonding pairs •• •• ••
Lone pairs • •• • •

37. Lewis Structures for Covalent Compounds

38. Lewis Structures
Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

39. Writing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
Keep track of the electrons:
(7) = 26

40. Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
Keep track of the electrons:
26 − 6 = 20

41. Writing Lewis Structures
Fill the octets of the outer atoms.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2

42. Writing Lewis Structures
Fill the octet of the central atom.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0

43. Writing Lewis Structures
If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.

44. Write the Lewis structure of CO2.

45. Writing Lewis Structures for Polyatomic Ions
The procedure is the same, the only difference is in counting the valence electrons.
For polyatomic cations, take away one electron from the total for each positive charge.
For polyatomic anions, add one electron to the total for each negative charge.

46. Example NO3─ 1. Write skeletal structure. 2. Count valence electrons.
N is central because it is the most metallic.
2. Count valence electrons.
N = 5
O3 = 3∙6 = 18
(-) = 1
Total = 24 e-

47. Tro's Introductory Chemistry, Chapter 10
Example NO3─ , Continued
3. Attach atoms with pairs of electrons and subtract from the total.
Electrons
Start 24
Used 6
Left 18
N = 5
O3 = 3∙6 = 18
(-) = 1
Total = 24 e-
Tro's Introductory Chemistry, Chapter 10 47

48. Example NO3─ , Continued 3. Complete octets, outside-in.
Keep going until all atoms have an octet or you run out of electrons.
N = 5
O3 = 3∙6 = 18
(-) = 1
Total = 24 e-
Electrons
Start 24
Used 6
Left 18
Electrons
Start 18
Used 18
Left 0

49. Example NO3─ , Continued
5. If central atom does not have octet, bring in electron pairs from outside atoms to share.
Follow common bonding patterns if possible.

50. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
ions or molecules with an odd number of electrons,
ions or molecules with less than an octet,
ions or molecules with more than eight valence electrons (an expanded octet).

51. Odd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

52. Fewer Than Eight Electrons
Consider BF3:
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.

53. Fewer Than Eight Electrons
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

54. Fewer Than Eight Electrons
The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

55. More Than Eight Electrons
The only way PCl5 can exist is if phosphorus has 10 electrons around it.
It is allowed to expand the octet of atoms on the third row or below.
Presumably d orbitals in these atoms participate in bonding.

56. Exceptions to the Octet Rule
H and Li, lose one electron to form cation.
Li now has electron configuration like He.
H can also share or gain one electron to have configuration like He.
Be shares two electrons to form two single bonds.
B shares three electrons to form three single bonds.
Expanded octets for elements in Period 3 or below.
Using empty valence d orbitals.
Some molecules have odd numbers of electrons. NO

57. Resonance

58. This is the Lewis structure we would draw for ozone, O3.
Resonance
This is the Lewis structure we would draw for ozone, O3. + -

59. Resonance
But this is at odds with the true, observed structure of ozone, in which…
…both O—O bonds are the same length.
…both outer oxygens have a charge of 1/2.

60. Resonance
One Lewis structure cannot accurately depict a molecule such as ozone.
We use multiple structures, resonance structures, to describe the molecule.

61. Drawing Resonance Structures
Draw first Lewis structure that maximizes octets.
Move electron pairs from outside atoms to share with central atoms.
If central atoms, 2nd row, only move in electrons, you can move out electron pairs from multiple bonds.

62. Practice—Draw Lewis Resonance Structures for CO32- (C Is Central with O Attached)62

63. Molecular Geometry

64. Molecular Shapes
The shape of a molecule plays an important role in its reactivity.
By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

65. What Determines the Shape of a Molecule?
Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

66. This molecule has four electron domains.
We can refer to the electron pairs as electron domains.
In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.
This molecule has four electron domains.

67. Valence Shell Electron Pair Repulsion Theory (VSEPR)
“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

68. Electron-Domain Geometries
These are the electron-domain geometries for two through six electron domains around a central atom.

69. Electron-Domain Geometries
All one must do is count the number of electron domains in the Lewis structure.
The geometry will be that which corresponds to that number of electron domains.

70. Molecular Geometries
The electron-domain geometry is often not the shape of the molecule, however.
The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

71. Molecular Geometries
Within each electron domain, then, there might be more than one molecular geometry.

72. Linear Electron Domain
In this domain, there is only one molecular geometry: linear.
NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

73. Trigonal Planar Electron Domain
There are two molecular geometries:
Trigonal planar, if all the electron domains are bonding
Bent, if one of the domains is a nonbonding pair.

74. Tetrahedral Electron Domain
There are three molecular geometries:
Tetrahedral, if all are bonding pairs
Trigonal pyramidal if one is a nonbonding pair
Bent if there are two nonbonding pairs

75. Trigonal Bipyramidal Electron Domain
There are four distinct molecular geometries in this domain:
Trigonal bipyramidal
Seesaw
T-shaped
Linear

76. Octahedral Electron Domain
All positions are equivalent in the octahedral domain.
There are three molecular geometries:
Octahedral
Square pyramidal
Square planar

77. Practice—Predict the Shape Around the Central Atom
PCl3

78. Practice—Predict the Shape Around the Central Atom
SiF5-

79. Molecular Polarity 79

80. Polarity of Molecules In order for a molecule to be polar it must:
1. Have polar bonds.
Electronegativity difference—theory.
Bond dipole moments—measured.
2. Have an asymmetrical shape.
Vector addition.
Polarity effects the intermolecular forces of attraction.

81. Molecule Polarity, Continued
The H—O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule. 81

82. Molecule Polarity
The O—C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule. 82

83. Example—Determining if a Molecule Is Polar83

84. Determine if NH3 is polar.

85. Example: Determine if NH3 is polar.
Write down the given quantity and its units.
Given: NH3
Build the Work area in this order: (1) “Write down ...”, (2) “Given:”.
Bold text “SeO2 ” in the Example box, duplicate it, and float to the Given line.
Make NEXT button hot.
When NEXT button clicked, float the text from the Given lines in the Work area to the upper right Information Box, then clear the Work area.

86. Example: Determine if NH3 is polar.
Information:
Given: NH3
Write down the quantity to find and/or its units.
Find: If polar

87. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Design a solution map.
Formula of compound
Molecular polarity
Lewis structure
Molecular shape
Bond polarity
and

88. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Draw the Lewis structure.
Write skeletal structure.

89. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Draw the Lewis structure.
Count valence electrons.
N = 5
H = 3 ∙ 1
Total NH3 = 8

90. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Draw the Lewis structure.
Attach atoms.
N = 5
H = 3 ∙ 1
Total NH3 = 8
Start 8 e-
Used 6 e-
Left 2 e-

91. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Draw the Lewis structure.
Complete octets.
N = 5
H = 3 ∙ 1
Total NH3 = 8 ∙∙
Start 2 e-
Used 2 e-
Left 0 e-

92. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Determine if bonds are polar.
Electronegativity
N = 3.0
H = 2.1 ∙∙
3.0 – 2.1 = 0.9
 polar covalent

93. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Determine shape of molecule.
4 areas of electrons
around N;
3 bonding areas
1 lone pair ∙∙
Shape = trigonal pyramid

94. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity
Apply the solution map.
Determine molecular polarity.
Bonds = polar
Shape = trigonal pyramid
Molecule = polar

95. Example: Determine if NH3 is polar.
Information:
Given: NH3
Find: If polar
Solution Map:
formula → Lewis → polarity and shape → molecule polarity ∙∙
Check:
Bonds = polar
Shape = trigonal pyramid
N = 5
H = 3 ∙ 1
Total NH3 = 8
The Lewis structure
is correct. The bonds
are polar and the
shape is unsymmetrical,
so the molecule should be polar.
Bonding = 3 ∙ 2 e-
Lone pairs = 1 ∙ 2 e-
Total NH3 = 8 e-
Molecule = polar

96. Practice—Decide Whether Each of the Following Molecules Is PolarEN
O = 3.5
N = 3.0
Cl = 3.0
S = 2.5 96

97. Practice—Decide Whether the Each of the Following Molecules Is Polar, Continued
Trigonal
bent Cl N O
3.0
3.5
Trigonal
planar O S
3.5
2.5
1. Polar bonds, N—O
2. Asymmetrical shape
1. Polar bonds, all S—O
2. Symmetrical shape
Polar
Nonpolar 97

98. Molecular Polarity Affects Solubility in Water
Polar molecules are attracted to other polar molecules.
Since water is a polar molecule, other polar molecules dissolve well in water.
And ionic compounds as well. 98

99. Properties of Substances
Ionic Bond

100. Properties of Substances
Metallic Bond

101. Properties of Substances
Covalent Bond

102. Molecular Solids Bond (Covalent) Nonmetal + Nonmetal Example
Ice Water, H2O
Melting Point
Low to moderate
Conductivity
Solid = no
Liquid = no
Aqueous = no
Malleability
brittle
Water, H2O

103. Metallic Solids Bond (Metallic) Example Copper, Cu Melting Point
Metal + Metal
Example
Copper, Cu
Melting Point
Low to High
Conductivity
Solid = yes
Liquid = yes
Aqueous = yes
Malleability
malleable
Copper metal, Cu

104. Ionic Solids Bond (Ionic) Example Table Salt, NaCl Melting Point High
Metal + Nonmetal
Example
Table Salt, NaCl
Melting Point
High
Conductivity
Solid = no
Liquid = yes
Aqueous = yes
Malleability
brittle
Table Salt, NaCl

105. Covalent Network Solids
Bond (Covalent)
Nonmetal + Nonmetal
Example
Diamond, C
Melting Point
Very High
Conductivity
Solid = no
Liquid = no
Aqueous = no
Malleability
Brittle, very hard