Section 6.5 – Molecular Geometry
Transcript of Section 6.5 – Molecular Geometry
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Section 6.5 – Molecular Geometry
• The properties of
molecules depend on
the bonding and the
molecular geometry,
the 3-dimensional
arrangement of the
atoms in space.
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Molecular Geometry
There are two
different theories
(VSEPR and
hybridization) both
based on
experimental
evidence, that can be
used to describe this.
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There Is No Conflict
• VSEPR Theory:
Accounts for molecular
bond angles.
• Hybridization:
Describes the orbitals
that contain the valence
electrons of a
molecule’s atoms.
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VSEPR Theory
Valence
Shell
Electron
Pair
Repulsion
Repulsion between the
valence-shell electrons
surrounding an atom causes
these sets to be oriented as
far apart as possible.
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Diatomic Molecules
• Diatomic molecules are
composed of two
atoms, so the geometry
is always linear, with
the molecular polarity
being determined by the
electronegativity
differences between the
atoms.
H2: non-polar HCl: polar
Electron distribution in N2, triple
bond in center of molecule, lone
pairs on the outside of each N.
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Shorthand for Describing• For molecules containing
more than two atoms, we can use the following symbols with subscripts:
A – the central atom
B – number of bonds
on the central atom
E – number of lone
pairs on central atom
(for atoms that have multiple bonds, it is treated as a single B for geometry)
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The Basis For VSEPR Theory
• That one must consider the locations of allelectron pairs of the valence electrons in the molecule. Polyatomic ions are treated the same way. The following do not have lone pairs that influence the geometry of the molecule.
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Linear – AB2
• Central atom with
two single bonds, no
lone pairs. Because
the valence electron
pairs in the bonds
repel each other, the
bonds are as far apart
as possible (180°).
• Ex: BeH2
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Trigonal Planar – AB3
• The 3 A-B bonds stay furthest apart by pointing to the corners of an equilateral triangle, giving 120°angles between the bonds.
Ex.: BH3
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Tetrahedral – AB4Octet rule is followed
here. The distance
between the A-B
bonds is maximized if
each bond points to the
corners of a
tetrahedron, giving
bond angles of 109.5°
between the bonds.
Ex.: CH4
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Trigonal-bipyramidal – AB5
• 120° angles
between bonds
within the
trigonal plane,
90° bond
angles between
the axial bond
and those in the
plane.Ex.: PCl5
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Octahedral – AB6
6 bonds to
the central
atom, all
equidistant
from each
other. 90°
bond
angles.
Ex.: SF6
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Lone Pairs Do Occupy Space and Influence
Geometry
• But our description of the molecular geometry refers to the positions of the atoms only.
• Because lone pairs on central atoms are still electrons (regardless of whether they are being shared), they repel other electrons around the central atom
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VSEPR and Unshared Electron Pairs
• One must always write
out the Lewis
structure for a
molecule to decide on
the proper geometry,
the chemical formula
of something does not
tell you about lone
pairs around the
central atom.
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Bent – AB2E
• 2 bonds to central atom
with one lone pair. The
lone pair bends what
one would expect to be
linear. The lone pair
takes up more space
than a bond and shoves
the bonded atoms closer
together than the 120°
for trigonal planar.
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Bent – AB2E2
• The addition of a
second lone pair forces
the bonding atoms even
closer together than
what one expects from
tetrahedral.
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Trigonal Pyramidal – AB3E
• With 3 bonds one
expects trigonal
planar with 120°
between the bonds,
but the lone pair
bends the plane away
from the pair, forcing
the atoms closer
together.
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Molecular Polarity
• Reflecting on the
geometries, we can now see
why lone pair on the central
atom make a molecule
polar—it changes the
geometry of the molecule
and creates an uneven
electron “tug of war” across
the molecule