CH 10: Molecular Geometry

Renee Y. Becker

Valencia Community College

CHM 1045

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Molecular Shapes: VSEPR

• The approximate shape of molecules is given by Valence-

Shell Electron-Pair Repulsion (VSEPR).

• Step 01: Count the total electron groups.

• Step 02: Arrange electron groups to maximize separation.

• Groups are collections of bond pairs between two

atoms or a lone pair.

• Groups do not compete equally for space:Lone Pair > Triple Bond > Double Bond > Single Bond

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Molecular Shapes: VSEPR

• Two Electron Groups: Electron groups point in opposite directions.

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Molecular Shapes: VSEPR

• Three Electron Groups: Electron groups lie in the same plane and point to the corners of an equilateral triangle.

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Molecular Shapes: VSEPR

• Four Electron

Groups:

• Electron groups

point to the

corners of a

regular

tetrahedron.

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Molecular Shapes: VSEPR

Five Electron Groups: Electron groups point to the corners of a trigonal bipyramid.

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Molecular Shapes: VSEPR

• Six Electron

Groups:

Electron groups

point to the

corners of a

regular

octahedron.

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Molecular Shapes: VSEPR

Electron Groups Lone Pairs Bonds Geometry Examples

2 0 2 Linear BeCl2

3 0 3 Trigonal planar BF3

3 1 2 Bent SO2

4 0 4 Tetrahedral CH4

4 1 3 Trigonal pyramidal NH3

4 2 2 Bent H2O

5 0 5 Trigonal bipyramidal PCl5

5 1 4 See-saw SF4

5 2 3 T-Shaped ClF3

5 3 2 linear I3-

6 0 6 Octahedral SF6

6 1 5 Square pyramidal SbCl52-

6 2 4 Square planar XeF4

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Example 1: VSEPR

• Draw the Lewis electron-dot structure and predict

the shapes of the following molecules or ions:

O3 H3O+ XeF2

PF6– XeOF4 AlH4

–

BF4– SiCl4 ICl4– AlCl3

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Valence Bond Theory

1. Covalent bonds are formed by overlapping of atomic orbitals, each of which contains one electron of opposite spin.

2. Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms.

3. The greater the amount of orbital overlap, the stronger the bond.

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Valence Bond Theory

• Linus Pauling: Wave functions from s orbitals & p orbitals could be combined to form hybrid atomic orbitals.

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• sp hybrid:

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• sp2 hybrid:

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• sp2 hybrid (π bond):

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• sp3 hybrid:

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sp3d hybrid:

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• sp3d2 hybrid:

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Hybridization Easy Way

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Molecular Orbital Theory

• The molecular orbital (MO) model provides a better explanation of chemical and physical properties than the valence bond (VB) model.

– Atomic Orbital: Probability of finding the electron within a given region of space in an atom.

– Molecular Orbital: Probability of finding the electron within a given region of space in a molecule.

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Molecular Orbital Theory

• Additive combination of orbitals () is lower in energy than two isolated 1s orbitals and is called a bonding molecular orbital.

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Molecular Orbital Theory

• Subtractive combination of orbitals () is higher in energy than two isolated 1s orbitals and is called an antibonding molecular orbital.

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Molecular Orbital Theory

• Molecular Orbital Diagram for H2:

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Molecular Orbital Theory

• Molecular Orbital Diagrams for H2–

and He2:

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Molecular Orbital Theory

• Additive and subtractive combination of

p orbitals leads to the formation of both

sigma and pi orbitals.

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Molecular Orbital Theory

• Second-Row MO Energy Level Diagrams:

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Molecular Orbital Theory

• MO Diagrams Can Predict Magnetic Properties:

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Molecular Orbital Theory

• Bond Order is the number of electron pairs

shared between atoms.

• Bond Order is obtained by subtracting the

number of antibonding electrons from the

number of bonding electrons and dividing by 2.

BO = Bonding electrons – antibonding electrons 2

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Example 2: Molecular Orbital Theory

• The B2 and C2 molecules have MO

diagrams similar to N2. What MOs are

occupied in B2 and C2, and what is the

bond order in each? Would any of these

be paramagnetic?

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B2

B has 5 electrons

So B2 has 10 elec

Core electrons don’t count toward BO

*2p

*2p

2p

2p

*2s

2s

*1s

1s 29

C2

*2p

*2p

2p

2p

*2s

2s

*1s

1s

Carbon has 6 electrons so C2 has 12 electrons

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Dipole moment

• Based on electronegativity differences between atoms in a molecule

• The most electronegative atom is partially negative

• The less electronegative atom is partially positive

• The dipole moment is the average of all dipoles in the molecule

• Exception (C-H bonds do not have a dipole)

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• If a molecule has a dipole moment it is polar

• A compound could contain polar bonds but the molecule could be non-polar because there is no dipole moment

• Bond dipoles can cancel each other out

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C OO C OO

Example 3: Dipole moment

• Draw the dipole moment for the following molecules, are they polar?

HCl NH3 CHCl3

H2O SF6 CCl4

CO2 CH2Cl2

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