1. Bonding and Molecular Structure
CHAPTER 8
Bonding and Molecular Structure

2. Introduction
Bonds: Attractive forces that hold atoms together in compounds
Valence Electrons: the outermost electrons
-These e- are involved in bonding

3. Valence Electrons
Electrons are divided between core and valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5

4. Valence Electrons
-The number of valence electrons of a main group atom is the Group number
-For Groups IA-IVA, number of bonding (unpaired)
electrons is equal to the group number
-For Groups VA -VIIA, number of bonding (unpaired)
electrons is equal to 8 - group number

5. Valence Electrons
-Except for H (and sometimes atoms of the 3rd group and higher)
-The total number of valence electrons around a given atom in a molecule will be eight:
OCTET RULE
- (with the exception of hydrogen) atoms in molecules prefer to be surrounded by 8 electrons (or have 4 bonds = 8 electrons)

6. Lewis Dot Formulas of Atoms
IA IIA IIIA IVA VA VIA VIIA VIIIA
H He
Li Be B C N O F Ne

7. Ionic Bonding
An ion is an atom or a group of atoms possessing a net electrical charge
-cations: positive (+) ions
These atoms have lost 1 or more electrons
-anions: negative (-) ions
These atoms have gained 1 or more electrons

8. Formation of Ionic Compounds
Monatomic ions consist of one atom
Examples:
Na+, Ca2+, Al3+ - cations
Cl-, O2-, N3- -anions
Polyatomic ions contain more than one atom
NH4+ - cation
NO2-,CO32-, SO42- - anions

9. Formation of Ionic Compounds
General trend:
metals become isoelectronic with the preceding noble gas electron configuration
nonmetals become isoelectronic with the following noble gas electron configuration

10. Formation of Ionic Compounds
Reaction of Group IA Metals with Group VIIA Nonmetals

11. Formation of Ionic Compounds
1s 2s p
Li 
F   
These atoms form ions with these configurations.
Li+  same configuration as [He]
F- same configuration as [Ne]

12. Formation of Ionic Compounds
In general:
the reaction of IA metals and VIIA nonmetals:
2 M(s) + X2  2 MX(s)
where M is the metals Li to Cs
and X is the nonmetals F to I
Electronically it looks like:
ns np ns np
M   M+ __ __ __ __
X    X-   

13. Formation of Ionic Compounds
reaction of IIA metals with VIIA nonmetals:
Be(s) + F2(g) BeF2(g)

14. Formation of Ionic Compounds
The valence electrons in these two elements react like:
2s p s p
Be [He]  __ __ __  Be2+ __ __ __ __
F [He]     F-   
Lewis dot structure representation:

15. Formation of Ionic Compounds
The remainder of the IIA metals and VIIA nonmetals react similarly:
M(s) X2  MX2
M can be any of the metals Be to Ba
X can be any of the nonmetals F to I

16. Formation of Ionic Compounds
For the reaction of IA metals with VIA nonmetals:
Draw the valence electronic configurations for Li, O, and
their appropriate ions

17. Formation of Ionic Compounds
Draw the electronic configurations for Li, O, and their appropriate ions
You do it!
2s p s p
Li [He]   Li1+
O [He]    O2-   
Draw the Lewis dot formula representation of this reaction

18. Formation of Ionic Compounds
Simple Binary Ionic Compounds Table
Reacting Groups General Formula Example
IA + VIIA MX NaF
IIA + VIIA MX BaCl2
IIIA + VIIA MX AlF3
IA + VIA M2X Na2O
IIA + VIA MX BaO
IIIA + VIA M2X Al2S3

19. Formation of Ionic Compounds
Reacting Groups General Formula Example
IA + VA M3X Na3N
IIA + VA M3X2 Mg3P2
IIIA + VA MX AlN
-H forms ionic compounds when bound to metals (IA and IIA metals
For example: LiH, KH, CaH2, and BaH2
-When H is bound to nonmetals, the compounds are covalent in nature

20. Formation of Covalent Bonds
-potential energy of an H2 molecule as a function of the distance between the two H atoms

21. Covalent Bonding Atoms share electrons If the atoms share:
2 electrons a single covalent bond is formed
4 electrons - a double covalent bond
6 electrons - a triple covalent bond
Atoms have a lower potential energy when bound…this is a more favorable situation (why?)

22. Writing Lewis Formulas:
1. Add the number of valence electrons for all the atoms that are present in the molecule
2. Add or subtract electrons based on the molecule’s (or ion’s) charge
3. Identify the central atom and draw a skeletal structure:
-the one that requires the most e- to complete octet
-the less electronegative
4. Place a bond between each atom (1 bond = 2 e-)
5. Fill in octet of outer atoms first
6. Finish by completing the octet of central atom
– if you run out of e- then multiple bonds must be created between the central atom and atoms bound to it

23. Writing Lewis Formulas
octet rule: representative elements usually attain stable noble gas electron configurations (8 valence e-) in most compounds
You must distinguish the difference between:
-bonding electrons and nonbonding electrons
-shared (paired) and unshared (unpaired) electrons

24. Formation of Covalent Bonds
Lewis dot structures:
1. H2 molecule formation:
2. HCl molecule formation:

25. Lewis Structures Homonuclear diatomic molecules
1. Two atoms of the same element, H2:
2. Fluorine, F2:
Nitrogen, N2:

26. heteronuclear diatomic molecules
Lewis Structures
heteronuclear diatomic molecules
1. hydrogen fluoride, HF
2. hydrogen chloride, HCl
3. hydrogen bromide, HBr

27. Lewis Structures
Water, H2O
Ammonia molecule , NH3

28. Lewis Structures Polyatomic ions: ammonium ion NH4+
Notice that the N-atom in this molecule has eight electrons around them (H does not)

29. Writing Lewis Formulas
Sulfite ion, SO32-.

30. Double and even triple bonds are commonly observed for C, N, P, O, and S
H2CO
SO3
C2F4

31. Lewis Structures
Example: Write Lewis dot and dash formulas for sulfur trioxide, SO3

32. Resonance There are three possible structures for SO3:
-Two or more Lewis formulas are necessary to show the bonding in a molecule
-use equivalent resonance structures to show the molecule’s structure
-Double-headed arrows are used to indicate resonance formulas

33. Resonance Resonance is a flawed method of representing molecules
-There are no single or double bonds in SO3

34. Sulfur Dioxide, SO2 1. Central atom = 2. Valence electrons = ___
or ___ pairs
3. Write the Lewis structure
4. Form double bond so that S has an octet — but note that there are two ways of doing this.

35. Limitations of the Octet Rule
There are some molecules that violate the octet rule:
- Be
- Group IIIA
-Odd number of total electrons.
-Central element must have a share of more than 8 valence electrons to accommodate all of the substituents. (i.e. S and P)

36. Limitations of the Octet Rule
Example: Write Lewis formula for BBr3.

37. Sulfur Tetrafluoride, SF4
Central atom =
Valence electrons = ___ or ___ pairs.
Form sigma bonds and distribute electron pairs.
5 pairs around the S atom. A common occurrence outside the 2nd period.

38. Limitations of the Octet Rule
Example: Write dot structures for AsF5.

39. Formal Atomic Charges Atoms in molecules often bear a charge (+ or -)
The predominant resonance structure of a molecule is the one with charges on atoms as close to 0 as possible
Formal charge = Group number – 1/2 (# of bonding electrons) - (# of Lone electrons)
= Group number – (# of bonds)
– (# of Lone electrons)

40. Formal Charge CO2
. .
. .
. .
. .

41. Which is the most stable resonance form?
Formal Charge
Thiocyanate Ion, SCN- • S N C • S N C • S N C
Which is the most stable resonance form?

42. Theories of Covalent Bonding
Valence Shell Electron Pair Repulsion Theory
Commonly designated as VSEPR
Principal originator
R. J. Gillespie in the 1950’s
Valence Bond Theory (Chapter 9)
Involves the use of hybridized atomic orbitals
L. Pauling in the 1930’s & 40’s

43. VSEPR Theory
electron densities around the central atom are arranged as far apart as possible to minimize repulsions (why?)
Five basic molecular shapes:
Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

44. VSEPR Theory
Two regions of high electron density around the central atom.

45. VSEPR Theory
2. Three regions of high electron density around the central atom.

46. VSEPR Theory
3. Four regions of high electron density around the central atom.

47. VSEPR Theory
4. Five regions of high electron density around the central atom.

48. VSEPR Theory
5. Six regions of high electron density around the central atom.

49. Molecular geometry: arrangement of atoms around the central atom(s)
VSEPR Theory
Electronic geometry(family): locations of regions of electron density around the central atom(s)
Molecular geometry: arrangement of atoms around the central atom(s)
Electron pairs are not used in the molecular geometry determination

50. VSEPR Theory
Lone pairs (unshared pairs) of electrons require more volume than shared pairs
-there is an ordering of repulsions of lone electrons around central atom
Criteria for the ordering of the repulsions:
1. Lone pair to lone pair is the strongest repulsion.
2. Lone pair to bonding pair is intermediate repulsion.
3. Bonding pair to bonding pair is weakest repulsion.

51. Molecular Shapes and Bonding
Symbolism:
A = central atom
B = bonding pairs around central atom
U = lone pairs around central atom
For example:
AB3U designates that there are 3 bonding pairs and 1 lone pair around the central atom

52. Linear Electronic Geometry: AB2
Some examples of molecules with this geometry:
BeCl2, BeBr2, BeI2, HgCl2, CdCl2

53. Trigonal Planar Electronic Geometry: AB3
Some examples of molecules with this geometry are:
BF3, BCl3

54. Tetrahedral Electronic Geometry: AB4
Some examples of molecules with this geometry are:
CH4, CF4, CCl4, SiH4, SiF4

55. VSEPR Theory
An example of a molecule that has the same electronic and molecular geometries is methane (CH4)
-Electronic and molecular geometries are tetrahedral

56. Tetrahedral Electronic Geometry: AB4

57. Tetrahedral Electronic Geometry: AB3U
Some examples of molecules with this geometry are:
NH3, NF3, PH3, PCl3, AsH3
-trigonal pyramidal
-electronic and molecular geometries are different.
. .
. .
107.5°

58. Tetrahedral Electronic Geometry: AB2U2
Some examples of molecules with this geometry are:
H2O, OF2, H2S
-bent
-electronic and molecular geometries are different
104.5°

59. VSEPR Theory
An example of a molecule that has different electronic and molecular geometries is water (H2O)
-Electronic geometry is tetrahedral
-Molecular geometry is bent or angular

60. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3
Some examples of molecules with this geometry are:
PF5, AsF5, PCl5
axial
equatorial
axial

61. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3
If lone pairs are incorporated into the trigonal bipyramidal structure, there are three possible new shapes:
One lone pair - Seesaw shape
Two lone pairs - T-shape
Three lone pairs – linear
The lone pairs occupy equatorial positions first:
-they are 120o from each other
-90o from the axial positions
Results in decreased repulsions compared to lone pair in axial position
axial
equatorial
axial

62. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3
AB4U molecules have:
trigonal bipyramid electronic geometry
seesaw shaped molecular geometry
polar
One example of an AB4U molecule is
SF4

63. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

64. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3
AB3U2 molecules have:
1. trigonal bipyramid electronic geometry
T-shaped molecular geometry
polar
One example of an AB3U2 molecule is
IF3

65. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

66. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3
AB2U3 molecules have:
trigonal bipyramid electronic geometry
linear molecular geometry
nonpolar
One example of an AB3U2 molecule is
BrF2-

67. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

68. Octahedral Electronic Geometry: AB6, AB5U, and AB4U2
Some examples of molecules with this geometry are:
SF6, SeF6, SCl6, etc.

69. Octahedral Electronic Geometry: AB6, AB5U, and AB4U2
If lone pairs are incorporated into the octahedral structure, there are two possible new shapes:
One lone pair - square pyramidal
Two lone pairs - square planar
The lone pairs occupy any position because they are all 90o from all bonds positions:
-Additional lone pairs occupy the position 180º from the first set of lone pairs
-This results in decreased repulsions compared to lone pairs in the other positions

70. Octahedral Electronic Geometry: AB6, AB5U, and AB4U2
AB5U molecules have:
octahedral electronic geometry
Square pyramidal molecular geometry
polar.
One example of an AB4U molecule is
IF5

71. Octahedral Electronic Geometry: AB6, AB5U, and AB4U2
AB4U2 molecules have:
octahedral electronic geometry
square planar molecular geometry
and are nonpolar.
One example of an AB4U2 molecule is
XeF4

72. Polarity and Electronegativity Figure 8.11

73. Dipole Moments
For example, HF and HI:

74. Dipole Moments some “nonpolar molecules” that have polar bonds
Two conditions to be polar:
1. There must be at least one polar bond present or one lone pair of electrons
2. the molecule must be nonsymmetric
Examples: water, CF4, CO2, NH3, NH4+

75. Polar Molecules Molecular geometry affects molecular polarity
-they either cancel or reinforce each other
A B A
A B A
linear molecule
nonpolar
angular molecule
polar

76. Polar and Nonpolar Bonds
Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds
-Nonpolar covalent bonds have a symmetrical charge distribution (electron distribution)

77. Polar and Nonpolar Bonds
Polar covalent bonds: electrons are not shared equally
-they have different electronegativities
H F
Electronegativities:
Difference = 1.9 very polar bond

78. Polar and Nonpolar Bonds
Compare HF to HI:
H I
Electronegativities:
Difference = 0.4 slightly polar bond
more complicated geometries exist…

79. Bond Polarity Three molecules with polar covalent bonds:
-Each bond has one atom with a slight negative charge (-d)
-another with a slight positive charge (+d)

80. Polar or Nonpolar?
AB3 molecules: BF3, Cl2CO, and NH3

81. Polar or Nonpolar?
CO2 and H2O
Which one is polar?

82. CH4 … CCl4 Polar or Not?
Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”

83. Compounds Containing Double Bonds
Ethene or ethylene, C2H4, is the simplest organic compound containing a double bond.
-has a double bond to obey octet rule
Lewis Dot Formula

84. Double Bonds
What is the effect of bonding and structure on molecular properties?
s and p
Free rotation around C–C single bond
No rotation around C=C double bond

85. # of bonds between similar pairs of atoms
Bond Order
# of bonds between similar pairs of atoms
Double bond
Single bond
Acrylonitrile
Triple bond

86. Bond Order
Consider NO2-:
The N—O bond order = 1.5

87. Bond Order (a) bond strength (b) bond length
Bond order is proportional to two important bond properties:
(a) bond strength
(b) bond length
745 kJ
414 kJ
110 pm
123 pm

88. Bond Length
the distance between the nuclei of two bonded atoms

89. Bond length depends on size of bonded atoms
H—F
Bond Length
H—Cl
H—I
Bond distances measured in Angstrom units where 1 Å = 10-2 pm.

90. Bond length depends on bond order
Bond distances measured in Angstrom units where 1 Å = 10-2 pm.

91. Bond Strength Measure of the energy required to break a bond
See Table 9.10
BOND STRENGTH (kJ/mol)
H—H KJ
C—C KJ
C=C KJ
CC KJ
NN KJ
The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond.