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9
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Examples of ABx Molecules and Polyatomic Ions
AB2   BeCl2, SO2, H2O,
AB3   BF3, NH3, ClF3,
AB4   CCl4,

, SF4, XeF4,
AB5   PCl5, IF5, SbF5, BrF5
AB6   SF6, UF6,

*The basis of the VSEPR model is that electron pairs in the
valence shell of an atom repel one another
...

*Each of the ABx molecules we consider will have one of these 5 electron domains geometries…
2
...
Trigonal Planar
4
...
Trigonal Bipyramidal
6
...


No lone pairs Bond Angle: the angle between two adjacent A-B bonds
on central
• AB2 molecule: bond angle is 180°
atom

• AB3 molecule: bond angle is 120°
• AB4 molecule: bond angle is 109
...


Electron-Domain and Molecular Geometries of Molecules with Lone Pairs on
the Central Atom
Total Number
of Electron
Type of
Domains
Molecule
3
AB2

ElectronDomain
Geometry

Number of
Lone Pairs
1

Placement of
Lone Pairs

Molecular
Geometry

Example
SO2

Bent
Trigonal planar
4

AB3

1

NH3
Trigonal
pyramidal

Tetrahedral
4

AB2

2

H 2O
Bent

Tetrahedral
5

AB4

1

SF4

SeesawTrigonal
bipyramidal
5

AB3

shaped
2

Trigonal
bipyramidal

ClF3

Tshaped

T-

Trigonal
bipyramidal
5

AB2

shaped
3

Linear

Trigonal
bipyramidal
6

AB5

1

Square
pyramidal

Octahedral
6

AB4

BrF5

2

XeF4

Square planar
Octahedral

*When there are lone pairs on the central atom in a trigonal pyramid, the lone pairs preferentially
occupy equatorial positions because repulsion is greater when the angle between electron
domains is 90° or less
...
Draw the Lewis structure
2
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Determine the electron-domain geometry by applying the VSEPR model
4
...

*A lone pair on a central atom is attracted only to the nucleus of that atom
...

• As a result, the lone pair has more freedom to spread out and greater capacity to
repel other electron domains
...


*Presence of lone pair means bond angles are less than ideal
...
2: Molecular Geometry and Polarity
*Molecular polarity influences physical, chemical, and biological properties
...


Because the sum of the two vectors is not zero in
both the x and y directions, the molecule is polar
...


both the x and y directions, the molecule is
nonpolar
...


Because the individual bond dipoles
sum to zero, the trans isomer is
nonpolar
...


*All equivalent angles are ideal
...

*Double bonds repel more than single bonds so they are greater than ideal
...
3: Valence Bond Theory
Bond Lengths and Bond Enthalpies of H2, F2, and HF
Bond Length (Å) Bond Enthalpy (kJ/mol)
H2 0
...
4

F2

1
...
6

HF 0
...
2

Valence Bond Theory: atoms share electrons when an atomic orbital on one atom overlaps with
an atomic orbital on the other
• Each of the overlapping atomic orbitals must contain a single, unpaired
electron
...

• The nuclei of both atoms are attracted to the shared pair of electrons
...

*A singly occupied orbital will appear as a light color and a doubly occupied orbital will appear
as a darker version of that same color
...

• The formation of covalent bonds is exothermic
...

*In summary, the important features of valence bond theory are as follows:
• A bond forms when singly occupied atomic orbitals on two atoms overlap
• The two electrons shared in the region of orbital overlap must be of opposite
spins
• Formation of a bond results in a lower potential energy for the system

9
...

*Hybrid orbitals are another type of electron
domain
...

*When you hybridize you always
end up with the same
number of orbitals you
started with
...
We
can think of the large lobe of each sp hybrid

The mixing of Be’s 2s orbital with one of its 2p
orbitals, a process known as hybridization yields
two hybrid orbitals that are neither s nor p but
have some character of each
...

orbital as the result of a
constructive combination and the
small lobe of each as the result of

a destructive combination
...
The realistic hybrid orbital shapes
are shown first
...
(b) The 2s orbital and one of the 2p orbitals on Be combine to form
two sp hybrid orbitals
...
Like any
two electron domains, the hybrid orbitals on Be are 180° apart
...


*Superscript numbers in hybrid orbital notation are used to designate the number of atomic
orbitals that have undergone hybridization
...
5: Hybridization in Molecules Containing Multiple Bonds
Sigma (σ) Bond: a bond in which the shared electron density is concentrated directly along the
internuclear axis
Pi π Bond: bonds that form from the interaction of parallel p orbitals
*A sigma bond and a pi bond together constitute a double bond
...


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Molecular Orbital Theory: a theory that describes the orbitals in a molecule as bonding and
antibonding combinations of atomic orbitals
Molecular Orbitals: an orbital that results from the interaction of the atomic orbitals for the
bonding atoms
• Specific shapes
• Specific energies
• Can each accommodate a maximum of two electrons
• Two electrons residing in the same molecular orbital must have opposite
spins
• The number of molecular orbitals equals the number of atomic orbitals we
combine
Bonding Molecular Orbital: a molecular orbital that is lower in energy than the atomic orbitals
that combined to produce it; has an electron density that pulls the nuclei together
Antibonding Molecular Orbital: a molecular orbital that is higher in energy than the atomic
orbitals that combined to produce it; has an electron density that pulls the nuclei in opposite
directions
Sigma (σ) Molecular Orbitals: molecular orbitals that lie along the internuclear axis
Bond Order: a number based on the number of electrons in bonding and antibonding molecular

directions
Sigma (σ) Molecular Orbitals: molecular orbitals that lie along the internuclear axis
Bond Order: a number based on the number of electrons in bonding and antibonding molecular
orbitals that indicates, qualitatively, how stable a bond is
• The higher the bond order, the more stable the molecule

*Molecular orbital theory predicts that a molecule with a bond order of zero will not exist
...
(a)
Bonding and antibonding
molecular orbitals shown
separately
...


*Filling of molecular orbitals follows the same rules as the filling of atomic orbitals…
1
...
Each orbital can accommodate a maximum of two electrons with opposite
spins
3
...

*When atomic orbitals of different energies interact to form molecular orbitals, the lower-energy
atomic orbital contributes more to the bonding molecular orbital and the higher-energy atomic
orbital contributes more to the antibonding molecular orbital
...

• The antibonding molecular orbital more closely resembles the atomic orbital
of the less electronegative atom
...


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Lewis Theory
a
...
To make qualitative predictions about bond strength and bond
lengths
ii
...
Widely used by chemists
b
...
Two-dimensional (molecules are three-dimensional)
ii
...
Fails to explain why bonds form
2
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Strength:
i
...
Weakness:
i
...
Valence Bond Theory
a
...
Describes the formation of covalent bonds as the overlap of atomic
orbitals à Bonds form because the resulting molecule has a lower
potential energy than the original, isolated atom
b
...
Alone fails to explain the bonding many molecules such as BeCl2,
BF3, and CH4 in which the central atom in its ground state does not
have enough unpaired electrons to form the observed number of
bonds

3

4

have enough unpaired electrons to form the observed number of
bonds
4
...
Strength:
i
...
Understand the bonding and geometry of more molecules,
including BeCl2, BF3, and CH4
b
...
Fail to predict some of the important properties of molecules, such
as the paramagnetism of O2
5
...
Strength:
i
...
Weakness:
i

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