Chapter Ten

1
Chapter Ten

• Bonding Theory
• and Molecular Structure

2
Molecular Geometry

• is simply the shape of a molecule.
• Molecular geometry is described by the geometric
  figure formed when the atomic nuclei are imagined
  to be joined by the appropriate straight lines.
• Molecular geometry is found using the Lewis
  structure, but the Lewis structure itself does
  NOT necessarily represent the molecules shape.

A water molecule is angular or bent.
3
VSEPR

• Valence-Shell Electron-Pair Repulsion (VSEPR) is
  a simple method for determining geometry
• Basis pairs of valence electrons in bonded
  atoms repel one another.
• These mutual repulsions push electron pairs as
  far from one another as possible.

What will be this B-A-B angle, when the electron
pairs (bonds) are as far apart as they can get?
4
Electron-Group Geometries

• An electron group is any collection of valence
  electrons, localized in a region around a central
  atom, that repels other groups of valence
  electrons.
• The mutual repulsions among electron groups lead
  to an orientation of the groups that are called
  electron-group geometry.
• These geometries are based on the number of
  electron groups

5
A Balloon Analogy
Each electron group may be -an unshared pair of
electrons, or -a bond (single, double, triple
bonds are each counted as one electron group).
6
VSEPR Notation

• In the VSEPR notation used to describe molecular
  geometries, the central atom in a structure is
  denoted as A, terminal atoms as X, and the lone
  pairs of electrons as E.
• Example ClF3 is designated AX3E2. It has three
  groups (atoms in this case) around the Cl atom,
  and two lone pairs of electrons on the Cl (draw
  the Lewis structure to see).
• For structures with no lone pairs on the central
  atom (AXn), the molecular geometry is the same as
  the electron-group geometry.
• When there are lone pairs, the molecular geometry
  is derived from the electron-group geometry.
• In either case, the electron-group geometry is
  the tool we use to obtain the molecular geometry.

7
(No Transcript)
8
(No Transcript)
9
(No Transcript)
10
(No Transcript)
11
Structures With No Lone Pairs

• AX2 both the electron-group geometry and the
  molecular geometry for two electron groups is
  linear.
• AX3 these molecules have a trigonal planar
  geometry.
• AX4 these molecules have a tetrahedral geometry.
• AX5 these molecules have a trigonal bipyramidal
  geometry
• AX6 these molecules have an octahedral geometry.
• The AX5 and AX6 require an expanded valence shell
  and, therefore, the central atom is a
  third-period or higher element.

12
Geometries of Methane
13
Structures With Lone Pairs

• Electron groups on the central atom repel one
  another, whether they are shared pairs or lone
  pairs.

• However, the geometry of the molecule is found
  using the bonded atoms.

Lone pair
The species is bent or angular, with a bond angle
of 120
The three electron groups are 120 apart
NO2
14
Some Structures With Lone Pairs

• AX2E these molecules have an electron-group
  trigonal planar geometry, but a bent molecular
  geometry.
• AX2E2 these molecules have an electron-group
  tetrahedral geometry, but a bent molecular
  geometry.
• AX3E these molecules have an electron-group
  tetrahedral geometry, but a trigonal pyramidal
  molecular geometry.
• AX4E these molecules have an electron-group
  trigonal bipyramidal geometry, but a seesaw
  molecular geometry.
• AX4E2 these molecules have an electron-group
  octahedral geometry, but a square-planar
  molecular geometry.

15
Molecular Geometry of Water
16
Polar MoleculesAnd Dipole Moments

• A polar bond was discussed in Chapter 9 it is a
  bond with separate centers of positive and
  negative charge.
• A molecule with separate centers of positive and
  negative charge is a polar molecule.
• The dipole moment (m) of a molecule is the
  product of the magnitude of the charge (d) and
  the distance (d) that separates the centers of
  positive and negative charge.
• m dd
• A unit of dipole moment is the debye (D).
• One debye (D) is equal to 3.34 x 10-30 C m.

17
Polar Molecules In An Electric Field
An electric field causes polar molecules to line
up but has no effect on nonpolar molecules.
18
Bond Dipoles AndMolecular Dipoles

• A polar covalent bond has a bond dipole a
  separation of positive and negative charge
  centers in an individual bond.
• Bond dipoles have both a magnitude and a
  direction (they are vector quantities).
• A molecule can have polar bonds, but may be a
  nonpolar molecule IF the bond dipoles cancel.

19
Bond Dipoles AndMolecular Dipoles

• CO2 has polar bonds, but is a linear molecule
  the bond dipoles cancel and it has no net dipole
  moment (m 0 D)

No net dipole

• The water molecule has polar bonds also, but is
  an angular molecule.
• The bond dipoles do not cancel (m 1.84 D), so
  water is a polar molecule.

Net dipole
20
Molecular ShapesAnd Dipole Moments

• Molecular polarity can be predicted based on the
  following three-step approach
• Use electronegativity values to predict bond
  dipoles.
• Use the VSEPR method to predict the molecular
  shape.
• From the molecular shape, determine whether bond
  dipoles cancel to give a nonpolar molecule or
  combine to produce a resultant dipole moment for
  the molecule.
• Lone pair electrons can also make a contribution
  to dipole moments.

21
Atomic Orbital Overlap

• Valence Bond (VB) theory states that a covalent
  bond is formed when atomic orbitals (AOs)
  overlap.
• In the overlap region, electrons with opposing
  spins produce a high electron charge density.

Overlap region between nuclei has high electron
density

• In general, the more extensive the overlap
  between two orbitals, the stronger is the bond
  between two atoms.

22
Bonding In H2S
The measured bond angle in H2S is 92 good
agreement
Hydrogen atoms s-orbitals can overlap with the
two half-filled p- orbitals on sulfur.
23
Points of VB Theory

• Most of the electrons in a molecule remain in the
  same orbital locations that they occupied in the
  separated atoms.
• Bonding electrons are localized in the region of
  AO overlap.
• For AOs with directional lobes (such as
  p-orbitals), maximum overlap occurs when the AOs
  overlap end to end.
• VB theory is not without its problems

24
Hybridization Of Atomic Orbitals

• VB theory carbon should have only two bonds,
  and they should be about 90 apart.

• Reality carbon has four bonds, which (singly
  bonded) are about 109 apart.

• We can hybridize the four orbitals
  mathematically combine the wave functions for the
  2s orbital and the three 2p orbitals on carbon.
• The four AOs combine to form four new hybrid AOs.
• The four hybrid AOs are equivalent, and each has
  a single electron (Hunds rule).

Four equivalent hybrid orbitals can now form four
bonds
25
sp3 Hybridization

• Hybridizing an s-orbital with three p-orbitals
  gives rise to four hybrid orbitals called (what
  else??) sp3 orbitals.
• The number of hybrid orbitals is equal to the
  number of atomic orbitals combined.
• The four hybrid orbitals, being equivalent, are
  about 109 apart.
• In bonding, hybrid orbitals may overlap with
  either pure atomic orbitals or with other hybrid
  orbitals.

26
The sp3 Hybridization Scheme
Four AOs
form four new hybrid AOs.
27
Methane and Ammonia
Four sp3 hybrid orbitals tetrahedral Four
electron groups tetrahedral Coincidence?
Hardly
In ammonia, one of the hybrid orbitals (top)
contains the lone pair that is on the nitrogen
atom
In methane, each hybrid orbital is a bonding
orbital
28
sp2 Hybridization

• Three sp2 hybrid orbitals are formed from an
  s-orbital and two p-orbitals.
• The empty p-orbital remains unhybridized. It may
  be used in a multiple bond.
• The sp2 hybrid orbitals are in a plane, 120o
  apart.
• This distribution gives a trigonal planar
  molecular geometry, as predicted by VSEPR.

29
The sp2 Hybridization Scheme in Boron
A 2p orbital remains unhybridized.
Three AOs combine to form
three hybrid AOs
30
sp Hybridization

• Two sp hybrid orbitals are formed from an
  s-orbital and a p-orbital.
• Two empty p-orbitals remains unhybridized the
  p-orbitals may be used in a multiple bond.
• The sp hybrid orbitals are 180o apart.
• The geometry around the hybridized atom is
  linear, as predicted by VSEPR.

31
sp Hybridization in Be
Two unused p-orbitals
32
Hybrid Orbitals Involvingd Subshells

• This hybridization allows for expanded valence
  shell compounds.
• By hybridizing one s, three p, and one d-orbital,
  we get five sp3d hybrid orbitals.
• This hybridization scheme gives trigonal
  bipyramidal electron-group geometry.
• By hybridizing one s, three p, and two
  d-orbitals, we get five sp3d2 hybrid orbitals.
• This hybridization scheme gives octahedral
  geometry.

33
The sp3d and sp3d 2 Hybrid Orbitals
sp3d
sp3d 2
34
Predicting Hybridization Schemes

• In the absence of experimental evidence, probable
  hybridization schemes can be predicted
• Write a plausible Lewis structure for the
  molecule or ion.
• Use the VSEPR method to predict the
  electron-group geometry of the central atom.
• Select the hybridization scheme that corresponds
  to the VSEPR prediction.
• Describe the orbital overlap and molecular
  geometry.

35
Hybrid Orbitals and TheirGeometric Orientations
36
Hybrid Orbitals AndMultiple Covalent Bonds

• Covalent bonds formed by the end-to-end overlap
  of orbitals, regardless of orbital type, are
  called sigma (s) bonds.
• All single bonds are sigma bonds.
• A bond formed by parallel, or side-by-side,
  orbital overlap is called a pi (p) bond.
• A double bond is made up of one sigma bond and
  one pi bond.
• A triple bond is made up of one sigma bond and
  two pi bonds.

37
VB Theory for Ethylene, C2H4
s-bond overlap of s-orbital of hydrogen and sp2
hybrid orbital.
p-bond has two lobes (above and below plane), but
is one bond. Side overlap of 2p2p.
s-bond sp2 - sp2 overlap
38
Summary of VB theory of Ethylene
39
VB Theory Acetylene
s-bond s - sp overlap
Two p-bonds (above and below, and front and back)
from 2p2p overlap
s-bond sp - sp overlap
form a cylinder of p-electron density around the
two carbon atoms.
40
Geometric Isomerism

• Geometric isomers are isomers that differ only in
  the geometric arrangement of certain substituent
  groups.
• Two types of geometric isomers include
• cis substituent groups are on the same side
• trans substituent groups are on opposite sides
• cis- and trans- compounds are distinctly
  different in both physical and chemical
  properties.
• Usually formed across double bonds and in square
  planar compounds.

41
Geometric IsomerismIn 2-Butene
Groups are on opposite sides of double bond
Groups are on the same side of double bond
42
trans-dichlorobis(ethylenediamine) cobalt(III)
chloride
43
Molecular Orbitals

• An alternative scheme to VB theory uses molecular
  orbitals.
• Molecular orbitals (MOs) are mathematical
  descriptions of the regions in a molecule where
  there is a high probability of finding electrons.
• Atoms atomic orbitals molecules molecular
  orbitals.
• In MO theory, molecular orbitals are formed by
  the combination of atomic orbitals.

44
Types of Molecular Orbitals

• A bonding molecular orbital (s or p) is at a
  lower energy level than the separate atomic
  orbitals that form it.
• Bonding orbitals have a high electron
  probability, or electron charge density, between
  the nuclei.
• An antibonding molecular orbital (s or p) is
  at a higher energy level than the separate atomic
  orbitals.
• Antibonding orbitals place a high electron
  probability away from the region between the
  bonded atoms.

45
Molecular Orbitals and BondingIn the H2 Molecule
Electrons in bonding orbitals increase the
stability of the molecule (compared to the
individual atoms).
46
Second Period Homonuclear Diatomic Molecules

• For every two atomic orbitals that are combined,
  two molecular orbitals result. A total of six
  molecular orbitals are formed from the six 2p
  atomic orbitals.
• Of each pair of molecular orbitals, one is a
  bonding molecular orbital at a lower energy than
  the separate atomic orbitals, and one is an
  antibonding orbital at a higher energy.

47
Molecular Orbitals Formed ByCombining 2p Atomic
Orbitals
48
Molecular Orbital Diagrams
Electrons fill MOs in the same way that AOs are
filled lowest energy to highest energy.
49
Bonding In Benzene

• Benzene (C6H6) was discovered by Michael Faraday
  in 1825.
• In 1865, Kekulé proposed that benzene has a
  cyclic structure, with a hydrogen atom attached
  to each carbon atom. Alternating single and
  double bonds join the carbon atoms.

• Modern view there are two resonance hybrids of
  benzene.
• The pi-electrons are not localized between any
  particular carbon atoms, but are delocalized
  among all six carbon atoms.

50
The s-Bonding Framework
51
Complete Structure of Benzene
Sigma bond between carbon atoms
and below the plane of sigma bonds.
Donut-shaped pi-cloud above
52
Aromatic Compounds

• Many of the first benzene-like compounds
  discovered had pleasant odors and hence acquired
  the name aromatic.
• In modern chemistry, the term aromatic compound
  simply refers to a substance with a ring
  structure and with bonding characteristics and
  properties related to those of benzene.
• All organic compounds that are not aromatic are
  called aliphatic compounds.

53
Some Aromatic Compounds
54
(No Transcript)
55
Summary

• The VSEPR method is used to predict the shapes of
  molecules and polyatomic ions.
• If all electron-groups are bonding groups, the
  molecular geometry is the same as the
  electron-group geometry.
• A polar covalent bond has separate centers of
  positive and negative charge, creating a bond
  dipole.
• In the valence bond theory, a covalent bond is
  formed by the overlap of atomic orbitals of the
  bonded atoms in a region between the atomic
  nuclei.
• Hybridized orbitals include sp, sp2, sp3, sp3d,
  and sp3d2.

56
Summary (continued)

• Unhybridized p orbitals overlap in a side-by-side
  fashion to form p bonds.
• Single bonds are all hybridized s bonds, double
  bonds have one s bond and one p bond, and triple
  bonds have one s bond and two p bonds.
• In molecular orbital theory, atomic orbitals of
  separated atoms are combined into molecular
  orbitals.
• The benzene molecule is usually represented by
  its resonance hybrid.
• Benzene-like compounds are called aromatic
  compounds.