Transcript 6.6 The Shape and Behavior of Molecules

6.6 The Shape and Behavior of
Molecules
Converting Lewis Structures into Three Dimensions — VSEPR Theory
In short, while remaining attached to the central atom, these groups of electrons will
position themselves as far away from each other as possible. This is the
fundamental principle behind the valence shell electron pair repulsion (VSEPR)
theory, which chemists use whenever they convert Lewis structures into molecular
shapes.
Converting Lewis Structures into Three
Dimensions — VSEPR Theory
If the electron groups are bonding electrons, then the peripheral atoms they bind to the
central atom adopt that same arrangement and produce a molecular shape.
Electron groups and their repulsive effects ultimately determine where and how the
nuclei of the atoms in a molecule or polyatomic ion arrange themselves in threedimensional space. And it’s the resulting shapes of those species that we really care
about.
We will consider each of the five electron group arrangements separately. Let’s begin by
assuming that each electron group is a bonding group connecting the central atom (with
single or multiple bonds) to the peripheral atoms. We’ll then expand our discussion to
include lone-pair electron groups on the central atom.
As a general notation, the central atom is “A,” a peripheral or surrounding atom is “X,”
and a lone-pair of electrons on the central atom is “E.”
Two-Bonding Electron Groups:
AX2
Three-Electron Groups:
AX3 and AX2E
Four-Electron Groups:
AX4, AX3E, and AX2E2
Five-Electron Groups:
AX5, AX4E, AX3E2, and AX2E3
Six-Electron Groups:
AX6, AX5E, and AX4E2
Molecular Formulas to Molecular
Shapes
We are now in a position to combine the information from this and the previous section to
predict molecular shapes starting with a molecular formula. The steps will guide you
through this process:
Step 1: Beginning with the formula, determine the Lewis structure using the steps
outlined in Section 6.4.
Step 2: Consider the central atom in the completed Lewis structure. Note the number
of bonded atoms and lone pairs associated with that atom.
Step 3: Assign an AXmEn notation to the molecule or polyatomic ion. (Note any bond
angles affected by the presence of one or more lone pairs.)
Step 4: Refer to the appropriate electron group arrangement category given in the
tables above to determine the shape of the molecule.
From Polar Bonds to Polar Molecules
Intermolecular Forces
Various Types
of Intermolecular Forces
Dipole-Dipole Forces — Attractions Between Polar Molecules
Hydrogen Bonds — Special Dipole-Dipole Forces
Dispersion (London) Forces — A Growing Attraction
Valence Bond Theory
The Hybridization of Atomic Orbitals and the Concept of
Formal Charge
Valence bond theory describes the location of bonding and non-bonding, lone electrons in
quantum-mechanical orbitals, created by overlap of the standard s, p, d, and f orbitals we
learned about earlier in Chapters 5 and 6. Hybridization is a mathematical process in which
these standard atomic orbitals are combined to form new atomic orbitals called hybrid
orbitals.
3
sp
Hybridization
This is the most common type of
hybridization as it is found in molecules
whose central atom is surrounded by a stable
octet. Methane, CH4 is a classic example of
sp3 hybridization. Each hydrogen has one
valence electron existing in a 1s orbital,
while carbon has four valence electrons, two
existing in a 2s orbital, and two in two lobes
of a 2p orbital. Carbon’s orbital diagram
would convert as shown in Figure 6.6.14
(notice the 1s electrons are not affected).
3
sp d
and
3
2
sp d
Hybridization
Sigma (σ) and Pi (π) Bonding
sp and
2
sp
Hybridization
Double and triple bonds are commonly associated with carbon atoms and always
consist of one sigma (σ) and one or two pi (π) bonds. In terms of hybridization, the
entire region of orbital overlap is considered one hybrid orbital, even though it may be
a double or triple bond. For this reason, multiple bonds are often associated with sp or
sp2 hybridization.
See summary table on page 392
Formal Charge
Formal charge is an entirely fictitious charge that is used to defend the most probable
Lewis structure of a molecule when several different Lewis structure arrangements are
possible. The formal charge on an atom in a molecule is the charge each atom would have if
all bonding electrons were shared equally between the bonded atoms. There are a couple of
ways to calculate the formal charge of each atom.