Chapter 11 Covalent Bonding Theories

1
Chapter 11Covalent Bonding Theories
2
Scientific models are approximations or
simplifications of reality
VSEPR predicts molecular shape by assuming that
e- groups minimize repulsions, and therefore
occupy as much space as possible around a central
atom Does NOT explain How molecular shapes
arise from interactions of atomic
orbitals Knowledge of molecular shape doesnt
help explain magnetic and spectral properties
of molecules (only an understanding of
orbitals and energy levels can do
that)
More than 1 theory is needed to explain complex
phenomena, such as covalent bonding, and the
resulting molecular shapes and behaviors
Valence bond (VB) theory molecular shape due to
interactions of atomic orbitals, which results in
new hybrid orbitals (sigma pi bonding two
types of covalent bonds) Molecular orbital
(MO) theory deals with orbitals associate with
the whole molecule (molecular orbitals) to
explain the energy and behavior of a molecule
3
Section 11.1 Valence Bond (VB) Theory
Based on the quantum mechanical model of the
atom (Chapter 7) an atom has certain allowed
(discrete) quantities of energy due to the
allowed frequencies of an electron whose
behavior is wavelike and whose exact location is
impossible to know
VB Theory a covalent bond forms when orbitals
of two atoms overlap and the overlap region,
which is between the nuclei, is occupied by a
pair of electrons
VSEPR Theory
VB Theory
4
Section 11.1 Valence Bond (VB) Theory

• Central themes of VB Theory
• The space formed by the overlapping orbitals has
  a maximum capacity of two e-s
• that must have opposite spins
• (2) The greater the orbital overlap, the stronger
  (and more stable) the bond.
• (Reason Bond strength depends on attraction of
  the nuclei for the shared e-s)
• (3) When two atoms form a covalent bonds, there
  orbitals overlap to form a new hybrid
• orbitals (which are not the original s-, p-,
  d- or f-orbitals, but some different shape)

5
Section 11.1 Valence Bond (VB) Theory
The extent of orbital overlap depends on the
shapes and directions of the atomic orbitals
s-orbitals are spherical ? Orientation during
bonding does not matter
2 e-
6 e-
In a bond involving p-, d-, and f- orbitals, the
orbitals will be oriented in a direction that
maximizes overlap. (If oriented in a direction
that does not maximize overlap, the bond will
be weaker)
10 e-
F-orbitals 7 orbitals, 2 e- in each
14 e-
6
Section 11.1 Valence Bond (VB) Theory

• Two points about formation of hybrid orbitals
  (called hybridization)
• of hybrid orbitals formed of atomic
  orbitals mixed
• type of hybrid orbital formed depends on the
  types of atomic orbitals mixed
• 5 common types of hybridization sp,
  sp2, sp3, sp3d, sp3d2

7
Section 11.1 Valence Bond (VB) Theory
sp hybridization sp hybrid orbitals (mix 1 s
1 p orbital of the central atom 2 hybrid
orbitals)
Linear e- group arrangement
Example BeCl2
Linear shape means that bonding orbitals must be
oriented linearly.
8
Section 11.1 Valence Bond (VB) Theory
VB Theory says Mixing two nonequivalent
orbitals (one s and one p) around a central
atom results in two equivalent sp hybrid orbitals
that lie 180º apart
Be hybridization 1s2, 2s2, 2p6 - NO e-s in the
p-orbitals of Be before it bonds
The 2s2, 2p6 orbitals of Be form the hybrid
orbital.
9
Section 11.1 Valence Bond (VB) Theory
sp hybrid orbital shape one small and one large
lobe
The 2s2, 2p6 orbitals of Be form the hybrid
orbital.
10
Section 11.1 Valence Bond (VB) Theory
sp hybrid orbital orientation during bonding e-
density extended in bonding direction, which
minimizes repulsion between e- occupying other
orbitals of the atom
11
Section 11.1 Valence Bond (VB) Theory
sp2 hybridization sp2 hybrid orbitals (mix 1 s
2 p orbitals of the central atom 3 hybrid
orbitals)
trigonal planar e- group arrangement
Bonding orbitals Be 1s2, 2s2 Cl 1s2,
2s2, 2p6, 3s2, 3p5 (Note The 2s2 orbital of
Be shares 1 e- with each Cl 3p5 1 of the three
2 p orbitals)
12
Section 11.1 Valence Bond (VB) Theory
sp3 hybridization sp3 hybrid orbitals (mix 1 s
3 p orbitals of the central atom 4 hybrid
orbitals)
tetrahedral e- group arrangement
CH4
NH3
H2O
13
Section 11.1 Valence Bond (VB) Theory
sp3d hybridization sp3d hybrid orbitals (mix 1
s 3 p 1 d orbitals of the central atom 5
hybrid orbitals)
trigonal bipyramidal e- group arrangement
For elements of Period (Row) 3 and higher
d-orbitals are included (can break octet rule)
14
Section 11.1 Valence Bond (VB) Theory
sp3d2 hybridization sp3d2 hybrid orbitals (mix
1 s 3 p 2 d orbitals of the central atom 6
hybrid orbitals)
octahedral e- group arrangement
15
Section 11.1 Valence Bond (VB) Theory
Summary
16
Section 11.1 Valence Bond (VB) Theory
When neither VB nor VSEPR theory apply
In hydrides When Group 6A (and sometimes 5A)
elements are the central atom.
H2S VSEPR tetrahedral geometry (ideal
109.5º) VB theory 4 sp3 hybrid orbitals
formed Reality Bond angle is 92º.
p-orbitals are unhybridized. Real factors
influence molecular shape Bond length Atomic
size Electron-electron repulsions Long bonds
between H and S result in less e- crowding and,
therefore, less e- repulsion. So, do not need
hybrid orbitals to minimize repulsion.
17
Section 11.1 Valence Bond (VB) Theory
In-class problems 11.8, 11.10,
11.12 Optional Homework problems 11.7,
11.9, 11.11, 11.19
18
Section 11.2 Modes of Orbital Overlap (s and p
bonds)
The s bond End-to-end overlap
All single bonds are s bonds.
Highest e- density is along the bond axis.
19
Section 11.2 Modes of Orbital Overlap (s and p
bonds)
The p bond Side-to-side overlap
All double bonds 1 s bond 1 p bond
Two regions of e- density in a p 1 above and 1
below the s bond axis (holds 2 e-). Significance?
Explains how double ( triple) bonds form
without e- repulsion.
20
Section 11.2 Modes of Orbital Overlap (s and p
bonds)
Triple bonds 1 s bond 2 p bond
21
Section 11.2 Modes of Orbital Overlap (s and p
bonds)
Bond strength of single, double, and triple bonds
End-to-end overlap of s bonds is more extensive
than side-to-side p bond overlap.
s bonds are strong than p bonds.
Based on this, is this statement True or
False Double bonds in ethylene are twice as
strong as single bonds in ethane, and triple
bonds in acetylene are three times as strong as
single bonds in ethane.
Other factors lone pair repulsions, bond
polarities, etc affect overlap and strength
22
Section 11.2 Modes of Orbital Overlap (s and p
bonds)
A final note on s and p bonds Rotation of one
part of molecule around another
s bonds allow free rotation, p bonds do not.
Significance? It is the reason that cis- and
trans- forms of molecules exist distinctly,
rather than as resonance hybrids.
23
Section 11.2 Modes of Orbital Overlap (s and p
bonds)
Problems 11.20, 11.21 Optional Homework
Problem 11.23