CHM 123 Chapter 8
8.1 Molecular shapes and VSEPR theory
VSEPR theory proposes that the geometric arrangement of terminal atoms, or groups of atoms about a central atom in a covalent compound, or charged ion, is determined solely by the repulsions between electron pairs present in the valence shell of the central atom
In the valence-shell electron-pair repulsion theory (VSEPR), the electron groups around a central atom
• are arranged as far apart from each other as possible, WHY?
• have the least amount of repulsion of the negatively charged electrons.
• have a geometry around the central atom that determines molecular shape.
The following are the “parent” electronic structures upon which VSEPR is based. These structures show how to minimize the energy of the structure by placing 2, 3, 4, 5 or 6 electron groups (charge clouds) as far apart around a central atom as possible in three-dimensional space.
Electrons in bonds and in lone pairs can be thought of as “charge clouds” that repel one another and stay as far apart as possible, this causing molecules to assume specific shapes.
Working from an electron-dot structure, count the number of “charge clouds,” and then determine the molecular shape.
Five electrons clouds
Six electron clouds
The Effect of Lone Pairs
· lone pair groups “occupy more space” on the central atom
o because their electron density is exclusively on the central atom rather than shared like bonding electron groups
· relative sizes of repulsive force interactions is:
· Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair
· this effects the bond angles, making them smaller than expected
How many electron groups (charge clouds) are around the central atom in the following? Identify the geometry shape
SO2 PCl5
Total # of e- groups on central atom / “Parent” electronic geometry / # Bondedatoms / # Lone
pairs / Idealized molecular shape / Idealized bond angles
2 / Linear / 2 / 0 / Linear / 180o
3 / Trigonal Planar / 3 / 0 / Trigonal Planar / 120 o
3 / Trigonal Planar / 2 / 1 / Bent / 120 o
4 / Tetrahedral / 4 / 0 / Tetrahedral / 109.5 o
4 / Tetrahedral / 3 / 1 / Trigonal Pyramidal / 109.5 o
4 / Tetrahedral / 2 / 2 / Bent / 109.5 o
5 / Trigonal Bipyramidal / 5 / 0 / Trigonal Bipyramidal / 90 o, 120 o, 180 o
5 / Trigonal Bipyramidal / 4 / 1 / Seesaw / 90 o, 120 o, 180 o
5 / Trigonal Bipyramidal / 3 / 2 / T-shaped / 90 o, 180 o
5 / Trigonal Bipyramidal / 2 / 3 / Linear / 180 o
6 / Octahedral / 6 / 0 / Octahedral / 90 o, 180 o
6 / Octahedral / 5 / 1 / Square Pyramidal / 90 o, 180 o
6 / Octahedral / 4 / 2 / Square Planar / 90 o, 180 o
Real bond angles vs. Idealized bond angles
VSEPR predicts the idealized bond angle(s) by assuming that all electron groups take up the same amount of space. Since lone pairs are attracted to only one nucleus, they expand into space further than bonding pairs, which are attracted to two nuclei. As a result, real molecules that has lone pairs on the central atom often have bond angles that are slightly different than the idealized prediction.
Central atom without lone pairs has the same real bond angle as the idealized angle.
The exceptions to this are square planar shapes and linear (derived from trigonal bipyramidal electronic structure) shapes where the lone pairs offset one another, thus causing no deviation from ideality.
Examples: Name the shape and give the idealized bond angles for the following
Lewis structure / Molecular Shape / Idealized bond angle / Real bond angleShapes of Larger Molecules – no one particular central atom
Many molecules don’t have a “central” atom but many “central” atoms. These molecules don’t fit into the shape names that we’ve learned. However, we can give an approximated shape and bond angle to each “central” atom at a time.
Example: Give the approximate molecular shape at the numbered “central” atoms
Drawing 3-D structures
Examples: Draw 3-D structure for the following molecules
8.2 –8.3 Valence Bond Theory and Hybridization
When a covalent bond is formed, there is shared electron density between the nuclei of the bonded atom
The simultaneous attraction of the shared electron density for both nuclei holds the atoms together, forming a covalent bond
Valence Bond Theory: A quantum mechanical model which shows how electron pairs are shared in a covalent bond.
Bond forms between two atoms when the following conditions are met:
• Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin.
• Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms.
• The greater the amount of overlap, the stronger the bond.
The Phase of an Orbital
In some cases, atoms use “simple” atomic orbital (e.g., 1s, 2s, 2p, etc.) to form bonds. In other case, they use a “mixture” of simple atomic orbitals known as “hybrid” atomic orbitals.
Two special names for covalent bonds of Organic molecules
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Dang6
Dang6
Dang6
Hybrid orbital – orbitals are used to describe bonding that is obtained by taking combinations of atomic orbitals of the isolated atoms.
1. When forming hybrid orbitals, the number of hybrid orbitals formed equals the number of orbitals mathematically combined or “mixed”. For example, if an s orbital is combined with a p orbital the result is two “sp” hybrid orbitals.
2. Hybrid orbitals have orientations around the central atom that correspond to the electron-domain geometry predicted by the VSEPR Theory
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How can the bonding in CH4 be explained?
sp3 Hybridization of Carbon
- Mixing 1s and all 3p atomic orbital
- Four sp3 hybridized orbitals equal in size, energy and shape
- Responsible for sigma bond ( single bond)
- Tetrahedral shape
Consider methane (CH4)
The sp2 Hybridization of Carbon
- Mixing 1s and 2p atomic orbitals
- Three sp2 hybridized orbitals equal in size, energy and shape
- One 2p unhybridized orbital
- Responsible for σ bond ( single bond)
- One π bond (double bond)
- Trigonal planar shape
3-D representation of ethane (C2H4)
The sp Hybridization of Carbon
- Mixing 1s and 1p atomic orbitals
- Two sp hybridized orbitals equal in size, energy and shape
- Two 2p unhybridized orbital
- one σ bond ( single bond)
- Two π bond (triple bond)
- Linear
Consider ethyne (C2H2)
3-D representation of ethyne (C2H2)
Five types of hybrid are shown below
# of e- groups around central atom / Hybrid orbitals used / Orientation of Hybrid Orbitals2 / sp
3 / sp2
4 / sp3
5 / sp3d
6 / sp3d2
Bonding to O and N
Like Carbon, O and N can participate in single bond and multiple bonds compose of σ and π
*Note: the lone pair or non bonding e- pair occupies space just as bonded atom
Examples:
1. Predict the hybridization, geometry, and bond angle for the carbon, oxygen and nitrogen atoms in the following molecules
2. What hybrid orbitals would be expected for the central atom in each of the following?
a) SF2 b) BrF3 c) ClF4+
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