Chapter Ten

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Chapter Ten

• Bonding Theory and
• Molecular Structure

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Contents
1. Molecular Geometry (1) Preview (2) VSEPR
Theory 2. Polar Molecules and Dipole Moments
3. Valence Bond (VB) Theory (1) Atomic Orbital
(AO) Overlap (2) Hybridization of Atomic
Orbitals (3) Multiple Covalent Bond 4. Geometric
Isomerism
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5. Molecular Orbital (MO) Theory (1) Characterist
ics of MOs (2) 1st Period Homonuclear Diatomic
MOs (3) 2nd Period Homonuclear Diatomic
MOs 6. Bonding in Benzene 7. Aromatic Compounds
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• Molecular Geometry
• Preview
• Molecular geometry The shape of a molecule which
  is described by the geometric figure for the
  atomic nuclei joined by straight lines.
• Some space-filling models for examples

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• VSEPR Theory
• Definition
• VSEPR (Valence-Shell Electron-Pair Repulsion )
  An approaching method for the geometric shapes of
  individual molecule or polyatomic ion basis of
  repulsion among the valence electrons associated
  with a central atom.

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• VSEPR theory proceeding
• Write a plausible Lewis structure
• Determine VSEPR notation
• AXmEn
• A Central atoms
• X Terminal atoms
• E Lone pairs electrons
• H2O for example

AX2E2
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• Determine the electron-group geometry

• An electron group can be
• - either single bond or a multiple bond
• - a (resonance) hybrid bond
• - a lone pairs of electron
• - a unpaired single-electron
• Repulsion force in general
• LP vs. LP gt LP vs. BP gt BP vs. BP
• Remark Lone-Pairs (LP), Bonding-Pairs (BP)
• Angle for repulsion forces 90 gt 120 gt 180
• For central atom belong to third-period or higher
  element with VSEPR notation such as AX5, AX4E,
  AX3E2, AX6, AX5E, AX4E2 require an expanded
  valence shell such as 3d orbital.

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• Balloon analogy to electron-group geometries

2 electron groups linear 3 electron groups
trigonal planar 4 electron groups tetrahedral 5
electron groups trigonal bipyramidal 6 electron
groups octahedral
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• Determine the molecular geometry
• Electron-group geometry Describing the
  arrangement of valence electrons about central
  atom.
• Molecular geometry Describing the arrangement of
  bonded atoms about the central atom.
• Structures with no lone-pair electrons (AXn
  type), two geometries are identical.
• Structures with lone-pair electrons (AXnEm type)
  type), two geometries are different.

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For AXn type molecule
Example 10.1 Use the VSEPR method to predict the
shape of the nitrate ion.

• Solution
• NO3-, total of valence electrons 5 (3 x 6)
  1 24
• A plausible Lewis structure

ii) VSEPR notation AX3 (no lone-pair electrons)
iii) Electron-group geometry
trigonal planar
trigonal planar ? O-N-O 120
iv) Molecular geometry
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For AXnEm type molecule
Example 10.2 Use the VSEPR method to predict the
molecular geometry of XeF2.

• Solution
• XeF2, total of valence electrons 8 (2 x 7)
  22
• Lewis structure

ii) VSEPR notation AX2E3 (2 BP, 3 LP)
Trigonal bipyramidal
iii) Electron-group geometry
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v) Molecular geometry three possibilities
Structure (I) one 120 LPLP, two 90 LPLP
repulsions. Structure (II) one 180 LPLP, two
90 LPLP Structure (III) no 90 LPLP, three
120 LPLP Repulsions 90 gt 120 gt
180 Structure (III) has the lowest energy
configuration. Ans The molecular geometry of
XeF2 is linear ? F-Xe-F 180
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For more than one central atom
Example 10.3 Use the VSEPR method to predict the
molecular geometry of HNO3.

• Solution
• HNO3, total of valence electrons 1 5 (3 x
  6) 24
• Lewis structure

Two central atom O and N

• VSEPR notation
• For central O atom AX2E2
• For central N atom AX3

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iii) Electron-group geometry For central O atom
tetrahedral For central N atom trigonal planar
iv) Molecular geometry For central O atom
angular For central N atom trigonal planar
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3) VSEPR geometry summary
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• Card for VSEPR
• i) Without lone-pair electrons

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ii) With lone-pair electrons
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5) Bond angle of molecular geometry i) Without
lone-pair electron
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ii) With lone-pair electron
Without lone-pair For comparison only
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• Polar Molecules and Dipole Moments
• (1) Polar molecule vs. Non-polar molecule
• Polar molecule A molecule in which the molecular
  dipole is nonzero.
• Non-polar molecule A molecule in which the
  molecular dipole is zero.
• (2) Bond dipole vs. Molecular dipole
• Bond dipole A separation of positive and
  negative charge in an individual bond.
• Molecular dipole

• For diatomic molecule molecular dipole is
  identical to bond dipole
• For a molecule consisted by three or more atoms,
  molecular dipole is estimated by the vector sum
  of individual bond dipole moment (also called net
  dipole or resultant dipole).

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• (3) Dipole moment (µ) - The polarity estimation
• For bond dipole
• A diatomic molecule with a polar covalent bond
  (such as HCl) for example

µ dd d magnitude of the partial charge (unit
coulomb, C) d distance that separates positive
and negative charge (unit meter, m) Unit for µ
Debye (D) 1 D 3.34 x 1030 C m

• For molecular dipole
• If the molecule consisted by three or more
  atoms, the molecular dipole is estimated by the
  vector sum of individual bond dipole moment (See
  problems 90 and 92)

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• (4) Polarity prediction
• Use electronegativity values to predict bond
  polarity
• Use VSEPR derived molecular shape and vector sum
  of individual bond dipoles to predict molecular
  polarity

CO2 Molecular geometry linear (net) dipole
moment m 0 D Nonpolar molecule
neutral
H2O Molecular geometry bent (net) dipole moment
m 1.84 D Polar molecule
d-
d
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(5) Polar molecules in electric field
Electric field ON Oriented regularly
Electric field OFF Oriented randomly
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Example 10.4 Explain whether you expect the
following molecules to be polar or nonpolar. (a)
CCl4 (b) CHCl3

• Solution
• Bond dipoles in all the bonds in CCl4 and in
  CHCl3
• ii) Lewis structures

iv) For CCl4 No molecular dipole Nonpolar
molecule For CHCl3 Downward molecular dipole
Polar molecule
iii) Molecular geometries
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• Valence Bond (VB) Theory
• The theory describes that covalent bonds are
  formed when atomic orbitals on different atoms
  overlap
• (1) Atomic Orbitals (AOs) Overlap
• Bonding in H2 for example

• A covalent bond is formed by the pairing of two
  electrons with opposing spins in the region of
  overlap of atomic orbitals between two atoms.
• This overlap region has a high electron charge
  density

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• (2) Hybridization of Atomic Orbitals
• Why hybridized (mixed) orbitals is necessary?

• Ground-state electron configuration of C for
  example, it should form only 2 bonds

• Indeed, each C can form 4 bonds

• Orbitals hybridization can solve this problem.

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• Some definition
• (Orbital) hybridization
• A process in which two or more AOs that are
  similar in energy but not equivalent, are
  combined to form a set of equivalent (hybrid
  atomic) orbitals that are properly oriented to
  form bonds.
• Hybrid atomic orbitals
• The new AOs formed from the process of
  hybridization which are equivalent in energy and
  oriented properly for forming bonds.
• Number of hybrid orbitals Number of hybrid
  orbitals is equal to the number of atomic
  orbitals combined, also equal to number of s bond
  number of lone electron pairs

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• Types of hybridization
• i) sp3 hybridization (C for example)

one s orbital with three p orbitals gives four
sp3 hybrid orbitals
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Examples of sp3 hybridization (for central atom)
Ground- State AO
Ele- ment
Mole- cule
Hybrid AO
C
2s
sp3
2p
N
2s
2p
sp3
O
2s
sp3
2p
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ii) sp2 hybridization (B for example)
one s orbital with two p orbitals gives three sp2
hybrid orbitals
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Examples of sp2 hybridization (for central atom)
Ground- State AO
Ele- ment
Hybrid AO
Molecule
B
2s
2p
sp2
2p
C
2p
2s
2p
sp2
2p
2s
N
2p
sp2
O
2p
2s
2p
sp2
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iii) sp hybridization (Be for example)
one s orbital with one p orbitals gives two sp
hybrid orbitals
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Examples of sp hybridization (for central atom)
Ground- State AO
Ele- ment
Hybrid AO
Molecule
Be
2s
sp
2p
2p
C
2s
sp
2p
2p
N
2s
sp
2p
2p
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• iv) Expanded valence shell hybridization
• sp3d hybridization
• PCl5 for example, P is central atom

Ground-state AO
Hybrid AO
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• sp3d2 hybridization
• SF6 for example, S is central atom

Ground-state AO
Hybrid AO
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• Procedures for predicting hybridization type
• a) Draw the Lewis structure
• b) Use the VSEPR method to predict the
  electron-group geometry of the central atom.
• c) Select the hybridization type that corresponds
  to the VSEPR prediction.

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• Reminding about hybridization
• Hybridization is employed for central atom only.
• Hybrid orbital describes the electron-group
  geometry for central atom.
• Total number of hybrid orbitals equal to the
  number of atomic orbitals combined, also equal to
  number of s bond number of lone electron pairs
• Total number of hybridization obitals of a
  central atom 2 ? sp 3 ? sp2 4 ? sp3 5
  ? sp3d
• 6 ? sp3d2.
• Molecular geometry is described by the relative
  atomic position around central atom
• Hybrid orbitals may overlap with pure atomic
  orbitals or with other hybrid orbitals forming s
  bond.

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vii) Example for electron-pair geometry/hybridizat
ion types versusmolecular geometry
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Example 10.6 Describe a hybridization scheme for
the central atom, and sketch the molecular
geometry of the IF5 molecule.
Solution

• Lewis
• structure

• VSEPR notation
• AX5E

iii) Hybrid AO

• Electron-group geometry octahedral
• Molecular geometry square pyramidal

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• Multiple Covalent Bond
• Some definition
• s (sigma) bond The (first) covalent bond formed
  by end-to-end overlap of pure or hybridized AO
  between the bonded atoms
• s s, s p, p p (end-to-end), s hybrid AO,
  p hybrid AO, hybrid AO hybrid AO
• p (Pi) bond The second (and third, if present)
  bond in a multiple bond, results from
  side-by-side overlap of unhybridized p AOs
• p p (side-by-side)
• - Single bonds one s bond
• - Double bond one s bond and one p bond
• - Triple bond one s bond and two p bonds

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• VB theory of bonding in ethylene (H2CCH2)
• example of a double bond

• A p-bond has two lobes (above and below plane),
  but is one bond, side-by-side overlap of 2p2p

• All six atoms in C2H4 lie in the same plane

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3) VB theory of bonding in Acetylene
(HC?CH) example of a triple bond

• Two p-bonds from 2p2p overlap forming a cylinder
  of p-electron density around the two carbon atoms

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• Example 10.7
• Formic acid, HCOOH, is the simplest carboxylic
  acid.
• Predict a plausible molecular geometry for this
  molecule.
• Propose a hybridization scheme for the central
  atoms that is consistent with that geometry.
• Sketch a bonding scheme for the molecule.

Solution

• Lewis structure

iii) Bonding scheme


Molecular geometry C trigonal planar O angular

• Hybridization
• C sp2, O sp3

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4. Geometric Isomerism

• Geometric Isomers The isomers that differ only
  in the geometric arrangement of certain
  substituent groups. Usually formed across double
  bonds and in square planar compounds.
• Two types of geometric isomers
• - 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.
• Example

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• Molecular Orbital (MO) Theory
• Molecular orbital (MO) A mathematical
  description of the region in a molecule where
  there is a high probability of finding electrons.
• MOs are formed by the combination of atomic
  orbitals, MOs electron configuration scheme to
  molecules like AOs electron configuration scheme
  to atoms.

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• (1) Characteristics of MOs
• 1) Two AOs forming Two MOs

• One is a bonding orbital, a high electron
  probability in the region between the bonded
  atoms, at a lower energy than separate AOs,
  electrons in bonding orbital increase the
  stability of the molecule.
• Another one an antibonding orbital, a high
  electron probability away from the region between
  the bonded atoms, at a higher energy than
  separate AOs, electrons in antibonding orbital
  decrease the stability of the molecule.

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2) MOs formed by combining s type AOs two 1s AOs
for example
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3) MOs formed by combining p type AOs two set 2p
AOs for example

• s2p and s2p
• end-to-end overlap of AOs
• p2p and p2p side-by-side overlap of AOs

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4) Applications of MOs

• Predict the existence of molecule
• Predicting magnetic properties
• Estimate the bond order (also bond length and
  bond energy)
• Bond Order (BO)
• (S bonding e - S antibonding e)/2

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(2) 1st Period Homonuclear Diatomic MOs H2 and
He2 for example
s1s
s1s
s1s
s1s
AOs of He (two 1s AOs)
AOs of H (two 1s AOs)
MOs of He2
MOs of H2
BO (2-0)/2 1 H2 molecule does
exist Diamagnetic
BO (2-2)/2 0 He2 molecule does not exist
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(3) 2nd Period Homonuclear Diatomic
MOs 1) Relative energy levels of MOs and electron
configurations
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2) Summary of 2nd period homonuclear diatomic MOs
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Lewis structure For O2

• Experiment showed O2 is paramagnetic
• MO prove O2 have unpaired electrons

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• Bonding in Benzene

• Kekulé proposed that benzene (C6H6) has a cyclic
  structure, with a hydrogen atom attached to each
  carbon atom.
• Alternating single and double bonds joined the
  carbon atoms together.

Resonance hybrid
Resonance structures

• All three p bonds are delocalized across all six
  carbon atoms

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• VB theory for benzene bonding showed six sp2-sp2
  s bonds

• MO theory for benzene bonding showed three p bonds

MO for p bonding in benzene

• The average bonder order of carbon-to-carbon in
  benzene (63)/6 1.5

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7. Aromatic compounds

• Some Definitions
• Aromatic compounds The organic compounds that
  have ring structure and bonding characteristics
  related to benzene, usually have pleasant odors.
• Aliphatic compounds The organic compounds that
  are not aromatic are called aliphatic compounds.
• Halogenated benzenes The organic compounds with
  halogen atoms substituted for one or more of the
  hydrogen atoms on a benzene ring.
• Function group C6H5 called phenyl

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(2) Some representative aromatic compounds
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(3) More aromatic compounds introduced
1,2-Dichlorobenzene (o-Dichlorobenzene)
1,3-Dichlorobenzene (m-Dichlorobenzene)
1,4-Dichlorobenzene (p-Dichlorobenzene)
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End of Chapter 10