Chapter 9
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The shape of a molecule is described by
reporting the locations of its atoms.
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In the VSEPR model, we focus on one atom at a
time, examining the bonds it forms and any
lone pairs it may possess. In order to
understand and study the properties of
molecules, we must be able to recognize and
understand the orientation of atoms within a
molecule.
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The atoms in a molecule that are bonded to each
other share a pair of valence electrons. Some
molecules also have atoms with nonbonding
electron pairs. Electron pairs repel each other, and
they want to take up a position in space that will
minimize their interactions with other electron
pairs, bonded or non bonded.
The goal of the VSEPR model is to arrange the
electron pairs around the central atom so that
there is the least amount of repulsions among
them. This occurs when the electron pairs are as
far away from each other as possible.
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To report shapes more precisely, we give the
bond angles, which are the angles between
bonds visualized as straight lines joining the
atom centers.
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The VSEPR model does not distinguish
between a single and multiple bonds. A
multiple bond is treated just like other region
of high electron concentration.
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Suggest a shape for the ethyne molecule,
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To understand any molecule, one must first
complete a Lewis dot structure. It is then
possible to predict the molecular shape using
Two basic Principles:
1. The shapes of molecules are determined by
the repulsion between electron pairs in the
outer shell of the central atom. Both bond pairs
(electron pairs shared by two atoms) and lone
pairs (those located on a central atom but not
shared) must be considered.
2. Lone pairs repel more than bond pairs.
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The molecule BF3 has a dot symbol as shown
below. Here the B atom has 3 bonded pairs in
its outer shell. Minimizing the repulsion causes
this molecule to have a trigonal planar shape,
with the F atoms forming an equilateral
triangle about the B atom. The F-B-F bond
angles are all 120°, and all the atoms are in the
same plane. Central atoms with octet
configurations:
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The molecule NH3 has a dot symbol much like
that for BF3. However, there is a lone pair in
the outer shell of the central N atom.
In NH3 the N has 3 bond pairs and 1 lone
pair, (4 total pairs). The shape is called trigonal
pyramidal (approximately tetrahedral minus
one atom).
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3 groups of electrons (3 bonds around a
central atom, no lone pairs)
BF3 Trigonal Planar
Molecule
4 groups of electrons (4 bonds around a
central atom, no lone pairs)
CF4 Tetrahedral
Molecule
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4 groups of electrons (3 bonds around a
central atom, 1 lone pair)
NH3 Trigonal pyramid
4 groups of electrons (2 bonds around a
central atom, 2 lone pairs)
H2O Bent Molecule
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4 groups of electrons (1 bond around a central
atom, 3 lone pairs)
HF Linear Molecule
5 groups of electrons (5 bonds around a
central atom, no lone pairs)
PF5 Trigonal Bipyramid
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5 groups of electrons (4 bonds around a
central atom, 1 lone pair)
SF4 See-saw Molecule
5 groups of electrons (3 bonds around a
central atom, 2 lone pairs)
ClF3 T-shaped Molecule
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5 groups of electrons (2 bonds around a
central atom, 3 lone pairs)
XeF2 Linear Molecule
6 groups of electrons (6 bonds around a
central atom, no lone pairs)
SF6 Octahedral Molecule
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6 groups of electrons (5 bonds around a
central atom, 1 lone pair)
BrF5 Square Pyramid
6 groups of electrons
(4 bonds around a central
atom, 2 lone pairs)
XeF4 Square Planar
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A covalent bond in which the electron pair is
shared unequally has partial ionic character and is
called a polar covalent bond.
A polar covalent bond is a bond between two
atoms that have partial electric charges arising
from their difference in electronegativity. Partial
charges give rise to an electric dipole moment.
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A diatomic molecule is polar if its bond is
polar. A polyatomic molecule is polar if it has
polar bonds arranged in space in such a way
that their dipoles do not cancel.
Water is a polar molecule because of the way
the atoms bind in the molecule such that there
are excess electrons on the oxygen side and a
lack or excess of positive charges on the
Hydrogen side of the molecule.
Water is a polar molecule with positive charges
on one side and negative on the other.
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Examples of polar molecules of materials that
are gases under standard conditions are:
Ammonia (NH3), Sulfur Dioxide (SO2) and
Hydrogen Sulfide (H2S).
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Page 413
9.52,9.58,9.60,
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The strength of a bond between two atoms is
measured by the bond enthalpy . The bond
enthalpy typically increases as the order of the
bond increases, decreases as the number of
lone pairs on neighboring atoms increases, and
decreases as the atomic radius increases.
The greater the bond strength the harder it will
be to break the bond.
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The first step in predicting the strength of a
bond in polyatomic molecule is to identify the
atoms and the bond order. An electronegative
atom can pull electrons toward itself from more
distant parts of the molecule. So all the atoms
in the molecule experience a slight pull.
Because these variations in strength are not
very great, the average bond enthalpy ∆HB is a
guide to the strength of a bond in any
molecule.
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The average bond enthalpies can be used to
estimate reaction enthalpies and to predict the
stability of a molecule.
Decide whether the reaction :
CH3CH2I(g) +H2O- CH3CH2OH(g) +HI(g)
is endothermic or exothermic.
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A bond length is the distance between the centers
of two atoms joined by a chemical bond. Bond
lengths affect the overall size of a molecule.
Bond length is directly related to bond order,
when more electrons participate in bond formation
the bond will get shorter. Bond length is also
inversely related to bond strength and the bond
dissociation energy, as a stronger bond is also a
shorter bond. In a bond between two identical
atoms half the bond distance is equal to the
covalent radius.
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The covalent radius of an atom is the
contribution it makes to the length of the
covalent bond. Covalent radii are added
together to estimate the lengths of bonds in
molecules.
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Pi bonds involve the electrons in the leftover
p orbital for each carbon atom. Those p orbitals
are the electron clouds or orbitals that are
shown going up above and below each carbon
atom
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The distinction between a sigma bond and a pi
bond is shown below. The sigma bond has
orbital overlap directly between the two nuclei.
The pi bond has orbital overlap off to the sides
of the line joining the two nuclei.
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A single bond is a σ- bond.
A double bond is a σ bond plus one π bond.
A triple bond is a σ bond and two π bonds.
In valence bond theory a bond forms when
unpaired electrons in valence shell atomic orbitals
on two atoms pair. The atomic orbitals they occupy
overlap end to end to form σ bonds or side by side
to form π bonds.
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Valence Bond theory cannot account for
bonding in polyatomic molecules like CH₄.
The electronic configuration of carbon is
[He] 2s²2px¹ 2py¹.The 2s orbital is filled in the
ground state of the carbon atom. Consider
forming hybrids by mixing the 2s orbital with
the 2pz orbital.
The configuration of carbon would be [He]
2s¹2px¹ 2py¹2pz¹
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Without promotion , a carbon atom has only
two unpaired electrons and so can form only
two bonds ; after promotion it has four
unpaired electrons and can form four bonds.
Each bond releases energy as it forms. Despite
the energy cost of promoting the electron, the
overall energy of methane molecule is lower
than it would be if carbon formed only two
C-H bonds.
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The carbon in methane is now sp³ hybridised.
Four hydrogen should form 4 bonds, one with
the 2s orbital of carbon and three with the 2p
orbitals in carbon.
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The four half-filled orbitals have equal energy
and are called sp3 hybrid orbitals. The s has a
superscript of 1 and the p has a 3, thus
meaning that there are four sp3 orbitals, all
with the same amount of energy. The hydrogen
atoms can now bond to these orbitals. This is
how methane forms.
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Hybridization also occurs in boron atoms,
where it forms sp2 orbitals. One electron from
the 2s orbital gets promoted to an empty 2p
orbital, thus forming three sp2 hybrid orbitals.
.
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The sp hybridization, occurs in beryllium. In
this example, one of the two electrons in the 2s
orbital gets 'promoted' to an empty 2p orbital,
thus forming two sp hybrid orbitals.
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This usually occurs when the central atom is an
element in period 3 or later, it uses its d orbital to
expand the valence shell to accommodate more
than 4 electron pairs.
The atom now uses one d orbital in addition to
four s and p orbitals of the valence shell. The
resulting orbital is called as sp³d hybrid orbital.
This alignment is Trigonal Bipyramidal because
the five outer atoms form two pyramids (one on
top, the other upside down on the bottom) around
the central atom. In this model, all atoms are 90
degrees or 120 degrees apart and have the orbital
hybridization of "sp3d"
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sp3d2
The final alignment is octahedral, not because it
has eight outer atoms, it actually has six, but
because it forms eight 90 degree angles. The
orbital hybridization is "sp3d2
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Paramagnetic substance means that it is
attracted to a magnetic field. It is a property of
unpaired electrons which act as tiny magnets.
Paramagnetic substances tend to move into a
magnetic field
Most substances contain fully paired electrons
and are repelled by a magnetic field ;such
substances are diamagnetic.
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Lewis theory cannot account for electron
deficient compounds or the paramagnetism of
oxygen. In molecular orbital theory, electrons
occupy orbitals that spread throughout the
molecule. The Pauli exclusion principle allows
no more than two electrons to occupy each
molecular orbital.
Constructive interference produces MO's that places
the bonding electrons between the nuclei of the
bonding atoms. Electrons are simultaneously
attracted by both nuclei which lowers their energies.
The MO's produced are called bonding orbitals.
 Destructive interference produces MO's that places
the electrons away from the space between the
nuclei. The nuclei repel which makes the bond less
stable (i.e. higher energy). These MO's are called
antibonding orbitals.

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Consider two helium atoms approaching. Two
electrons go into the sigma 1s bonding MO,
and the next two into the sigma star
antibonding MO. Bond order=(#of electrons in
bonding orbital-#of electrons in antibonding
orbitals)/2
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As antibonding MOs are more antibonding
than bonding MOs are bonding, He2
(dihelium), is not expected to exist
And dihelium has a bond order of zero.
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