Covalent Bonding

1
Chapter 3

• Covalent Bonding

2
History of Bonding Theories

• Gilbert N. Lewis (1875-1946)

3
Lewis Shortcoming

• O2 has been experimentally determined to possess
  a double bond plus two unpaired electrons
  (paramagnetic)
• Electron dot diagrams cannnot be drawn to satisfy
  BOTH of these criteria

4
Molecular Orbital (MO) Theory

• the study of the resulting molecular orbitals
  from the overlap of atomic orbitals
• In covalent molecules, it is called the linear
  combination of atomic orbitals (LCAO theory)

5
MO Theory

• Electrons placed in bonding orbitals create a
  more stable energy

6
MO Theory

• For orbitals to overlap, the signs on the
  overlapping lobes must be the same

7
MO Theory

• Whenever two atomic orbitals mix, two molecular
  orbitals are formed. The bonding orbital is
  always lower in energy

8
MO Theory

• For significant mixing to occur, the atomic
  orbitals must be of similar energy

9
MO Theory

• Each molecular orbital can hold a maximum of two
  electrons, one with spin ½ , the other ½

10
MO Theory

• The electron configuration of a molecule can be
  constructed by using the Aufbau principle by
  filling the lowest energy molecular orbitals in
  sequence.
• Li2 ?1s2 ?1s2 ?2s2

11
MO Theory

• When electrons are placed in different molecular
  orbitals of equal energy, the parallel
  arrangement (Hunds Rule) will have the lowest
  energy

12
MO Theory

• The bond order in a diatomic molecule is defined
  as the number of bonding electrons minus the
  number of antibonding electrons divided by two.

13
First Row Diatomic MOs

• H2

14
First Row Diatomic MOs

• H2

15
First Row Diatomic MOs

• He2

16
First Row Diatomic MOs

• He2

17
Second Row Diatomic MOs

• Due to increased Zeff, the 1s orbitals are pulled
  closer to the nucleus and are not involved in
  bonding.
• The 2s and 2p are involved in bonding and are
  called frontier orbitals.

18
Second Row Diatomic MOs

• Li2 and Be2

Li2 (?1s)2(?1s)2(?2s)2 Be2
(?1s)2(?1s)2(?2s)2 (?2s)2
19
Mixing of p Orbitals

• End on end and side to side mixing

20
MO Configurations of Mixing of p Orbitals
21
Second Row Diatomics
22
Heteronuclear Diatomics

• Different atomic orbitals will have different
  energies
• the higher the Zeff, the lower the atomic orbital
  energy

23
Heteronuclear Diatomics

• CO
• The s orbital on the oxygen atom is too low in
  energy to interact with the s orbital on the
  carbon atom, so a non-bonding orbital is formed

24
Heteronuclear Diatomics

• CO
• Due to the rotation axis, only the pz orbital on
  the oxygen atom has the appropriate symmetry to
  interact with the s orbital on carbon

25
Heteronuclear Diatomics

• CO
• The px and py orbitals on both carbon and oxygen
  are of the correct orientation to interact

26
Heteronuclear Diatomics

• CO
• The pz orbital is left unused on the carbon atom
  which remains as a nonbonding orbital

27
Heteronuclear Diatomics

• CO
• Bond order is correctly predicted as 3 (6
  bonding electrons 0 antibonding electrons)/2

28
Heteronuclear Diatomics

• HCl

29
Lewis Theory

• Most simplistic view of bonding
• Useful for deducing molecular shape
• Octet rule

30
Lewis Theory

• 4 rules to constructing electron-dot diagrams
• identify the central atom (lower
  electronegativity), place the symbols of the
  other atoms around the central atom
• NF3

31
Lewis Theory

• 4 rules to constructing electron-dot diagrams
• count the total number of valence electrons
• NF3
• N He2s22p3 5 valence electrons
• F He2s22p5 7 valence electrons each
• 5 (3x7) 26 total valence electrons

32
Lewis Theory

• 4 rules to constructing electron-dot diagrams
• place an electron pair between the central atom
  and each of the surrounding atoms, add lone pairs
  to the surrounding atoms
• 26 total valence electrons

33
Lewis Theory

• 4 rules to constructing electron-dot diagrams
• If the number of electrons on the central atom is
  less than 8, and there are left-over electrons,
  add lone pairs to the central atom. If there are
  no more electrons, use lone pairs from
  surrounding atoms
• 26 total valence electrons

34
Exceptions to the Octet Rule

• Rule is only generally applicable to Period 2
  elements
• boron only has six, hydrogen only two
• many elements have more than eight

35
Partial Bond Order

• Resonance structures

36
Partial Bond Order

• Bond order of 1 1/3

37
Formal Charge
1.
2.
3.

• To find formal charge, divide the bonding
  electrons equally among the atoms, subtract the
  resulting electrons from the original number of
  valence electrons
• For 1, N (5-6) -1, N (5-4) 1, O (6-6) 0
• For 2, N (5-5) 0, N (5-4) 1, O (6-7) -1
• For 3, N (5-7) -2, N (5-4) 1. O (6-5) 1

38
Formal Charge
-1 1 0
0 1 -1
-2 1 1

• The lowest energy structure will have the
  smallest formal charges on the atoms
• Resulting bond order of NN of 2 ½ and NO of 1
  ½

39
VSEPR Theory

• Valence shell electrons will be located as far
  from each other as possible
• good for deducing molecular shape but not much
  else
• concerned with electron groupings around the
  central atom
• single, double, and triple bonds are all
  considered one grouping

40
VSEPR Theory

• Linear geometry
• two electron groups, no lone pairs
• 180 bond angles

41
VSEPR Theory

• Trigonal planar geometry
• three electron groups, no lone pairs
• 120 bond angles

42
VSEPR Theory

• Bent (V-shaped or angular) geometry
• three electron groups, one lone pair
• 120 bond angles (many deviations)

43
VSEPR Theory

• Tetrahedral geometry
• four electron groups, no lone pairs
• 109.5 bond angles

44
VSEPR Theory

• Trigonal pyramidal geometry
• four electron groups, one lone pair
• 109.5 bond angles (many deviations)

45
VSEPR Theory

• Bent (V-shaped or angular) geometry
• four electron groups, two lone pairs
• 109.5 bond angles (many deviations)

46
VSEPR Theory

• Trigonal bipyramidal geometry
• five electron groups, no lone pairs
• 120 equatorial bond angles
• 90 axial bond angles

47
VSEPR Theory

• Seesaw geometry
• five electron groups, one lone pair
• 120 equatorial bond angles (many deviations)
• 90 axial bond angles (many deviations)

48
VSEPR Theory

• T-shaped geometry
• five electron groups, two lone pairs
• 120 equatorial bond angles (many deviations)
• 90 axial bond angles (many deviations)

49
VSEPR Theory

• Linear geometry
• five electron groups, three lone pairs
• 180 bond angles

50
VSEPR Theory

• Octahedral geometry
• six electron groups, no lone pairs
• 90 bond angles

51
VSEPR Theory

• Square-based pyramidal geometry
• six electron groups, one lone pair
• 90 bond angles (many deviations)

52
VSEPR Theory

• Square planar geometry
• six electron groups, two lone pairs
• 90 bond angles

53
VSEPR Theory

• Structures with seven electron groups
• pentagonal bipyramidal
• capped trigonal prismatic
• capped octahedral

54
Valence-Bond Theory

• Quantum mechanical adaptation of the Lewis theory
  by Pauling
• Summarized in four statements
• A covalent bond results from the pairing of
  unpaired electrons in neighboring atoms
• The spins of the paired electrons must be
  antiparallel
• To provide enough unpaired electrons in each atom
  for the maximum bond formation, electrons can be
  excited to fill empty orbitals
• The shape of the molecule results from the
  directions in which the orbitals of the central
  atom point

55
Valence-bond Theory

• Valence-bond theory originally did not account
  for certain bond angles
• NH3
• should point in the x, y, and z directions
  according to bonding using nitrogens px, py, and
  pz orbitals (90 angles)
• observed angles are 107

56
Orbital Hybridization

• Wavefunctions of electrons in atomic orbitals can
  mix together to form hybrid orbitals
• sp, sp2, sp3, sp3d, and sp3d2 are common
  hybridizations
• number formed is the sum of the atomic orbitals

57
Orbital Hybridization
Orbitals Orbitals Orbitals Type of Hybridization Number of Hybrid Orbitals Resulting Molecular Geometry
s p d Type of Hybridization Number of Hybrid Orbitals Resulting Molecular Geometry
1 1 0 sp 2 Linear
1 2 0 sp2 3 Trigonal Planar
1 3 0 sp3 4 Tetrahedral
1 3 1 sp3d 5 Trigonal Bipyramidal
1 3 2 sp3d2 6 Octahedral
58
Orbital Hybridization

• BF3
• boron has an electron configuration of
  He2s22p1, gets excited to He2s12p2
• orbitals mix to provide three sp2 hybridized
  orbitals at 120 angles

59
Orbital Hybridization

• CO2
• carbon has an electron configuration of
  He2s22p2, gets excited to He2s12p3
• the s and one p orbital mix to provide two sp
  hybridized orbitals at 180 angles
• single electrons in the two remaining p orbitals
  interact with a p orbital on oxygen to give a ?
  bond

60
Network Covalent Substances

• A crystal lattice which consists solely of
  covalently bound atoms
• diamond and quartz
• high melting points, hard, and insoluble

61
Intermolecular Forces

• If no intermolecular forces existed, all
  substances would be gases
• All molecules exhibit London Dispersion Forces
• induced dipole attractions
• Other forces are dipole-dipole, ion-dipole, and
  hydrogen bonding

62
London (Dispersion) Forces

• Separation of charges lead to a temporary dipole
• this leads to a partial positive and negative
  charge which will attract themselves between
  molecules

63
London (Dispersion) Forces

• Strength of the dispersion forces is related to
  the number of electrons and shape
• The higher the force, the higher the melting and
  boiling points

64
Electronegativity

• Power of an atom to attract electrons
• greater for higher Zeff
• dipoles can cancel each other out as in CO2

65
Dipole-Dipole Forces

• Attractions between permanent dipoles
• Induced dipole attractions are much higher than
  dipole-dipole forces

66
Hydrogen Bonding

• Exceptionally strong dipole-dipole forces
• strongest intermolecular force
• similar to a weak, covalent bond