Molecular Shapes

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Dr. Saidane
Chem. 152
Molecular Shapes
Skills you should have mastered
Conceptual
1.
2.
3.
4.
Distinguish between single, double and triple bonds.
Define a polyatomic ion.
Define resonance.
Explain the significance of electronegativity and what the ionic or covalent
character of a bond means.
5. Explain how the polarity of a bond depends on the electronegativities of the two
atoms in the bond.
6. Relate the properties of a molecule to its shape.
7. Explain the VSEPR theory.
8. Explain how the shapes of molecules affect their polarity.
9. Distinguish between polar bond and polar molecule.
Problem-solving
1.
2.
3.
4.
5.
Write Lewis structures for an ionic, metallic, and covalent compound.
Determine the formula of an ionic compound using the Lewis structure and the
crisscross method.
Write Lewis structures and resonance forms of polyatomic ions and molecular
compounds.
Use a table of electronegativities to predict the polarity of bonds.
Using the VSEPR theory and the Lewis structure, predict the shape, the bond
angles, and the polarity of molecules with single bonds, multiple bonds and
lone pairs.
Lewis Structures for Covalent Compounds
Nonmetals share electrons with one another until they have completed their octet (or duplet for
hydrogen). A Lewis structure of a covalent compound shows the arrangements of electrons as
shared pairs (or lines) and lone pairs (pairs of dots).
To write the Lewis structure of a covalent compound we need to follow these steps:
1.
Calculate the total number of valence electrons in the compound. Ex: CH4 has 8 valence
electrons (4 electrons for C and one electron for each H).
2. Arrange the atoms in the compound. Choose the central atom and place the remaining
atoms symmetrically around the central atom. Ex: SO2 is OSO not SOO.
3. Place one electron pair (or a line) between each pair of bonded atoms. Each bonded pair of
atoms must have at least a single bond between them.
4. Complete the octet of each atom (duplet for H) by placing the remaining valence electron
as electron pairs around the atoms. If there are not enough electron pairs form multiple
bonds.
Exception to the Octet Rule
Some elements do not have an octet in their Lewis structures. Hydrogen and boron are the
elements than can have a valence shell with less than eight electrons. H has two electrons
(duplet). Boron tends to form bonds in which six electrons surround it, such as in BF3.
Resonance Structures
In some Lewis structures of covalent compounds with multiple bonds, the multiple bonds can
be written in several equivalent locations. For each Lewis structure of the compound with a
different location of the multiple bond we have a resonance structure and each one of them is called
resonance hybrid.
SHAPES OF MOLECULES AND IONS
To a chemist, one of the most interesting aspects of a molecule is its shape or
molecular geometry, which is the arrangement in space of the atoms bonded to each
other. Lewis structures do not directly show the three-dimensional arrangements of
atoms in a molecule, which play a key role in determining chemical reactivity. For
example, C3H8 have different Lewis structures called isomers.
Each isomer has
different physical properties (smell, boiling point, etc.) and chemical properties (react
differently when placed in the presence of the same compound). To account for these
molecular shapes, we need just one addition to Lewis’s model: the VSEPR model.
The VSEPR Model
The VSEPR model is the Valence-Shell Electron-Pair Repulsion Model.
This
model is based on the fact that regions of high electron concentration repel one
another. Bonding electrons and lone pairs take up positions as far from one another as
possible, to minimize their repulsion and therefore maximize their separations.
In this model, we focus on the central atom of a molecule and we imagine that all the
electrons involved in bonds to this atom and the lone pairs belonging to the central
atom lie on the surface of an invisible sphere that surrounds the central atom. These
bonding electrons and lone pairs are regions of high electron concentration, and they
repel one another.
To minimize these repulsions, these regions move as far as
possible on the surface of the sphere. We identify the shape of the molecule by noting
where the atoms lie and look up the name of the shape.
Multiple Bonds in the VSEPR Model
When using the VSEPR model to predict molecular shapes, we do not need to
distinguish between single and multiple bonds. A multiple bond is treated as another
region of high electron concentration just as a single bond.
Molecules With Lone Pairs on The Central Atom
If a molecule has lone pairs on the central atom, they occupy a region around the
central atom and therefore contribute to the shape of the molecule. However only the
positions of atoms are considered when describing the shape of a molecule, and
therefore lone pairs are ignored when we name the shape of the molecule.
The Distorting Effect Of Lone Pairs
According to the VSEPR model, lone pairs have a more strongly repelling effect than
bonding pairs. In other words it is best for lone pairs to be as far as possible from
each other as possible. It is also best for atoms to be far from lone pairs, even though
that might bring them close to other atoms. The strengths of repulsions are in the
order:
lone pair-lone pair > lone pair-bonding air > bonding pair-bonding-pair.
How To Use the VSEPR Model
The general procedure for predicting the shape of a molecule is as follows:
1. Write the Lewis structure of the molecule. Treat a multiple bond equivalent of a
single electron pair. Identify the central atom and assign the letter A to the
central atom.
2. Identify the number of bonds (bonding domain) and lone pairs (nonbonding
domain) around the central atom. Each lone pair will be assigned the letter E,
and each bond will be assigned the letter X.
3. Classify the molecule as an AXnEm, n is the number of bonds and m the number
of lone pairs around the central atom. Determine the shape according to the
class.
4. Allow the molecule to distort so that lone pairs are as far as possible from one
another and from bond pairs.
The shape of the molecule is described in the following table:
n
m
Shape
2
0
Linear
3
0
Trigonal Planar
2
1
Bent
4
0
Tetrahedral
3
1
Trigonal Pyramidal
2
2
Bent
CHARGE DISTRIBUTION IN MOLECULES
We need to know where electrons are most likely to be found in molecules if we want
to understand a molecule’s chemical properties.
Polar Bonds
A polar covalent bond is a bond between two atoms that share unequally the bonding
electrons arising from their difference in electronegativity. In this case an atom will
have a partial positive charge and the other atom in the bond will have a partial
negative charge. Partial charges give rise to electric dipole moment. The size of an
electric dipole is a measure of the magnitude of the partial charges.
Polar Molecules
A polar molecule is a molecule with a non-zero dipole moment. A nonpolar molecule
is a molecule that has zero electric dipole moment.
The shape of a molecule governs whether or not it is polar. Highly symmetric
arrangements of polar bonds result in molecules that are nonpolar, because in these
molecules the bond dipoles cancel.
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