1. Bonding General Concepts
Chapter 8
Bonding General Concepts

2. Types of Bonding Ionic Bonding Covalent Bonding Metallic Bonding
Occurs when atoms gain or lose electrons to become ions
Very Strong Attractions
Covalent Bonding
Occurs when atom share electrons
Metallic Bonding
Occurs when metal atoms allow a “sea of electrons” to be shared

3. Examples Ionic Covalent Metallic
Sodium chloride, Lithium Sulfate, Iron (II) Chloride
Covalent
Carbon dioxide, Octane, Ethanol
Metallic
Aluminum, Copper, Bronze

4. But How Do I Know Ionic Covalent Metallic Metals and Nonmetals
Metals and Polyatomic Ions
Polyatomic Ions and Polyatomic Ions
Covalent
Nonmetals
Metallic
Metals

5. Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
Developed by Linus Pauling
Values range 0.7 to 4.0
Fluorine = 4.0
Francium and Cesium = 0.7
What is the periodic trend?
Top to Bottom – Decrease
Left to Right Increase

7. H----F Electron Sharing
Electrons are not always shared evenly in covalent bonds.
Called Polar Covalent Bonds
Example of HF
H----F

10. But How Do I Know Revisited
Ionic
Between atoms with a large difference in electronegativity
Nonpolar Covalent
Between atoms with no difference in electronegativity
Polar Covalent
Between atoms with a medium difference in electronegativity

11. Ion Size Ions are not the same size as their parent atom
Positive Ions are smaller than parent
Negative Ions are larger
For a group of isoelectronic ions the most positive ion is the smallest

13. Example
Place the following ions in order of decreasing size
Na+, K+, Rb+, Cs+
Cs+> Rb+> K+ > Na+
Se-2, Br-, Rb+, Sr+2
Se-2> Br-> Rb+> Sr+2

14. Ionic Compound Formation
The formation of ionic compounds from their elements is an exothermic process
Several energy aspects that must considered

15. Energy Considerations
What must happen for the reaction
Na(s) + 1/2Cl2(g)  NaCl(s)
We need to get Na+ and Cl- ions
Sublimation of Na
Ionization of Na (Ionization Energy)
Breaking Cl2 bond (Bond Energy)
Ionizing Cl (Electron Affinity)
Combination of Na+ and Cl- (Lattice Energy)
Sum is ΔHfº = Energy Change

16. Example
#42 p. 404
Find ΔHfº
Mg(s) + F2(g)  MgF2(s)
Lattice Energy = kJ/mol
Sublimation of Mg = 150 kJ/mol
First Ionization Energy = 735 kJ/mol
Second Ionization Energy = 1445 kJ/mol
Bond Energy = 154 kJ/mol
Electron Affinity = -328 kJ/mol

17. Lattice Energy Comparisons
The lattice energy for two sets of ions can be compared with a form of Coulomb’s Law
K is a constant (don’t worry about it)
Q1 and Q2 are the charges of the ions
R is the distance between the centers of the ions
The LE will be neg. if the charges are opposite

18. Example
Compare the lattice energies of Sodium Fluoride and Magnesium Oxide
Sodium Fluoride will have a smaller LE because of the smaller charges

21. Homework
p. 403 #’s 21,22,23,30,35abce, 40,41,46

22. Lewis Structures
A method for determining the arrangement of bonds in covalent species
Similar to dot structures but shows all bonds present

23. How 2 Draw Determine the number of valence electrons
Determine the central atom.
Usually the single atom, or the one in the middle of the formula
Place other atoms around the middle and bond
Complete octets with remaining electrons
If each atom does not have 8 electrons
Multiple Bonds my be necessary

24. Bond Types Single Bonds = 2 electrons Double Bonds = 4 electrons
Weakest and longest bonds
Double Bonds = 4 electrons
In the middle
Triple Bonds = 6 electrons
Strongest and shortest
There are not quadruple bonds!

25. General Rules Hydrogen will only form single bonds
Halogens usually only form 1 bond. Why?
7 valence electrons
Oxygen will have 2 bonds and often forms multiple bonds
Carbon likes to form chains

26. Example – Draw the Lewis Structure for CBr4

27. Example – Draw the Lewis Structure for NF3

28. Example – Draw the Lewis Structure for O2

29. Example – Draw the Lewis Structure for CS2

30. Example – Draw the Lewis Structure for BeF2

31. Example – Draw the Lewis Structure for C6H14

32. Polyatomic Ions
Atoms that are covalently bonded together and have a charge
Lewis structure rules
Negatively charged add electrons
Positively charged subtract electrons
Place Lewis structure in brackets when you are finished

33. Example – Draw the Lewis Structure for NO+

34. Resonance Species where equivalent Lewis structures exist
Electron density is spread out evenly between resonant bonds
Delocalized – Spread out
Often present in polyatomic ions

35. Example – Draw the Lewis Structure for CO3-2

36. Example – Draw the Lewis Structure for NO2-

37. Example – Draw the Lewis Structure for AsF5

38. Homework
p. 405 #’s 61,63,65,72

39. Formal Charge
Difference between the number of valence electrons on free atom and the valence electrons in a species
FC=Valence Electrons on free atom – valence electrons on the species
Atoms desire lowest formal charge possible
Negative formal charge should reside with most electronegative element

40. Example
Use formal charge to compare the molecules

41. Molecular Geometry
Lewis Structures do not show us the shape of molecules
Use VSEPR Theory
Valence Shell Electron Pair Repulsion Theory
Electron Groups want to be as far apart as possible in molecules
1 Electron Group = Single, Double or Triple Bond or Lone Pair of Electrons
Lone Pair Decrease the Bond Angle

42. 2 Electron Groups
Name Linear
Bond Angle 180

43. 3 Electron Groups
Name Trigonal Planar
Bond Angle 120

44. 3 Electron Groups 1 Lone Pair
Name Bent
Bond Angle <120

45. 4 Electron Groups
Name Tetrahedral
Bond Angle 109.5

46. 4 Electron Groups 1 Lone Pair
Name Trigonal Pyramidal
Bond Angle 107

47. 4 Electron Groups 2 Lone Pairs
Name Bent
Bond Angle 104.5

48. 5 Electron Groups
Name Trigonal Bipyramidal
Bond Angle 90 &120

49. 5 Electron Groups 1 Lone Pair
Name SeeSaw or Distorted Tetrahedral

50. 5 Electron Groups 2 Lone Pairs
Name T Shape

51. 5 Electron Groups 3 Lone Pairs
Name Linear

52. 6 Electron Groups
Name Octrahedral
Bond Angle 90

53. 6 Electron Groups 1 Lone Pair
Name Square Pyramidal

54. 6 Electron Groups 2 Lone Pairs
Name Square Planar

55. 6 Electron Groups 3 Lone Pairs
Name T Shape

56. 6 Electron Groups 4 Lone Pairs
Name Linear

57. Molecular Polarity
A polar molecule is one that has a partially positive and partially negative side
Molecules are Always nonpolar if they are one of the 5 base shapes w/ the same atom at the ends
Molecules are Always polar their bond dipoles do not cancel out
Molecules are polar if they do not have the same atoms at the end

58. Bond Dipole
Example the Bond Dipole for CO2 and CH2O

59. Predict whether the molecule is polar or nonpolar

60. Homework
P. 406 #’s 73,77,82,86 explain,92

61. Bonding Carbon forms four bonds with Hydrogen but, how!
Carbon [He] 2s p2
There are only 2 electrons to share
Something more must have to happen!

62. Hybridization
Mixing of different energy orbitals to form new bonding orbitals
In CH4 Carbon needs to blend 1 s orbital and 3 p orbitals to be able to bond
Called sp3 hybridization
4 electron groups gives sp3 hybridization

65. What is a Bond? σ (Think of it as a single bond)
A bond is the overlap of orbitals
Two hybrid orbitals, a hybrid and a nonhybrid, or two nonhybrid
First bond to form is called a sigma bond
σ (Think of it as a single bond)

66. Draw C2H6 in terms of orbitals
How are the H’s aligned?

67. sp2 Hybridization Blending of 1s and 2p orbitals
Used for 3 electron group geometry
There is still 1 unhybridized p orbital left over
Runs perpendicular to hybrid orbitals
Unhybridized p is used for double bond
Called a pi bond (π)

71. Draw C2F4 in terms of orbitals
How are the F’s aligned?
How many sigma bonds? Pi bonds?

72. sp Hybridization Blending of 1 s and 1 p orbital
Used for 2 electron group geometry
There is still 2 unhybridized p orbitals left over
Run at 90 degrees of each other

74. Draw HCN in terms of orbitals
How many sigma bonds? Pi bonds?

75. Expanded Octets
Some atoms can expand their octets by utilizing unused d orbitals
Must be in period 3 or greater
5 electron groups uses 1 d orbtital
dsp3 hybridized
6 electron groups uses 2 d orbitals
d2sd3 hybridized

76. Draw PCl5 in terms of orbitals

77. Draw SF6 in terms of orbitals

78. Bond Order The number of bonds between two atoms Ex. H2 is 1 O2 is 2
N2 is 3

79. Homework
P. 441 #’s 11-15,22,24,28d