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Unit 9 - Molecular Geometry, vsepr-geometries, image - Coggle Diagram
Unit 9 - Molecular Geometry
9.1 - Molecular Shapes
Lewis Structures show bonding and lone pairs, but do not denote shape.
The overall shape of a molecule is determined by its bond angles
Central atom A us bonded to n B atoms (ABn)
Five basic geometric structures of the molecule
Linear = 180° between two surrounding atoms
Trigonal planar = 120° between three surrounding atoms
Tetrahedral = 109.5° between four surrounding atoms
Trigonal bipyramidal = 90° and 120° between five surrounding atoms
Octahedral = 90° between six surrounding atoms
9.2 - The VSEPR Model
The Valence-Shell Electron-Pair Repulsion (VSEPR) Model allows us to predict shapes of molecules based on their bonding and nonbonding electron pairs.
From these patterns, bond angles between elements can be predicted along with the polarity of the molecules.
We can refer to the directions to which electrons point as electron domains. This is true whether there is one or more electron pairs pointing in that direction.
To determine the electron-domain geometry, count the total number of lone pairs, single, double, and triple bonds on the central atom.
Once you have determined the electron-domain geometry, use the arrangement of the bonded atoms to determine the molecular geometry.
Remember that some elements can break the octet rule and make more than four bonds (or have more than four electron domains).
For larger molecules, look at the geometry about each atom rather than the molecule as a whole. The larger molecule can have multiple molecular geometry types.
9.3 - Molecular Shape and Molecular Polarity
The polarity of bonds is based on the electronegativity difference between two atoms which explains how equally (or unequally) the electrons are shared between the atoms.
Molecules can also be described as polar and nonpolar. However, the polarity of a molecule depends on both the polarity of the bonds but also the symmetry of the molecular geometry.
To determine if a molecule is polar or nonpolar ask yourself:
Are the bonds COVALENT or IONIC?
If IONIC, the molecule is polar.
If COVALENT:
Are the BONDS polar?
a.NO: The molecule is NONPOLAR!
b.YES: Continue—Do the AVERAGE position of δ+ and δ– coincide?
1)YES: The molecule is NONPOLAR.
2)NO: The molecule is POLAR.
9.4- Covalent Bonding and Orbital Overlap
The VSEPR model is effective in predicting the shapes of molecules but does not explain why bonds exist.
The Valence-Bond theory helps explain why bonds exist by combining the Lewis electron diagrams with atomic orbital theory.
In Valence-Bond Theory, electrons of two atoms begin to occupy the same space.
This is called “overlap” of orbitals.
The sharing of space between two electrons of opposite spin results in a covalent bond.
Increased overlap brings the electrons and nuclei closer together until a balance is reached between the like charge repulsions and the electron-nucleus attraction.
Atoms can’t get too close because the internuclear repulsions get too great.
9.5 - Hybrid orbitals
Hybrid orbitals form by “mixing” of atomic orbitals to create new orbitals of equal energy, called degenerate orbitals.
When two orbitals “mix” they create two orbitals; when three orbitals mix, they create three orbitals; etc.
sp orbitals:
Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.
These sp hybrid orbitals have two lobes like a p orbital.
One of the lobes is larger and more rounded, as is the s orbital.
These two degenerate orbitals would align themselves 180° from each other.
sp2 and sp3 Hybrid Orbitals
Whenever we mix a certain number of atomic orbitals, we get the same
number of hybrid orbitals
Hypervalent Molecules
The elements which have more than an octet (exceptions to the octet rule)
Valence-Bond model would use d orbitals to make more than four bonds.
This view works for period 3 and below. Theoretical studies suggest that the energy needed would be too great for this
9.6- Multiple Bonds
Sigma (σ) bonds which are single bonds with the electrons concentrated along the line connecting the two nuclei.
Head to Head overlap
Cylindrical symmetry of electron density about the internuclear axis.
Pi (π) bonds occur in multiple bonds (double and triple) bonds, having a sideways overlap between two p orbitals because they are oriented perpendicular to the internuclear axis.
Side to Side overlap
Electron density above and below the internuclear axis
Weaker than σ bonds because there is less overlap
Single bonds are always σ-bonds.
Multiple bonds have one σ-bond, all other bonds are π-bonds.
