Transcript Molecular Structure and Covalent Bonding Theories

Molecular Structure and Covalent
Bonding Theories
Chapter 8
The Valence Shell Electrons
• Valence shell electrons
– These electrons are largely responsible for____
– Electrons not present in the preceding ___ ___
• Ignore filled sets of d and f orbitals
– Used to determine the Lewis structure of a
compound containing covalent bonds
• Works well for molecules containing atoms from the
____ ____ elements
Models to Describe Covalent
Bonding
• Valence shell electron pair repulsion
(VSEPR) model – predicts the _____
______ of atoms in a molecule
– This will be related to a physical property
called ______
• Valence bond theory – predicts how
bonding will take place by ______ of
atomic orbitals
VSEPR Theory
• Valence shell electron are present as either ___ ___ or
____ ____.
– Regions of high electron density are created.
– These regions arrange themselves to be as far away as
possible form on another. As a result specific geometries are
created around atoms in the molecule
• Single, double, and triple bonds are counted as one region of electron
density
• Unshared pairs of valence electrons are also counted as one region of
electron density
Drawing the Lewis structure accurately will reveal the number
of electron density regions around the center atoms
VSEPR Theory
• Draw the Lewis structures for CO2, H2CO,
and CH4
– How will these regions of electron density
arrange themselves to be as far away as
possible from one another?
• There are five basic shapes based on the
number of electron density regions around a
center atom(s)
– Illustration of models with next few slides
VSEPR Theory
Two regions of high electron density
VSEPR Theory
Three regions of high electron density
VSEPR Theory
Four regions of high electron density
VSEPR Theory
Five regions of high electron density
VSEPR Theory
Six regions of high electron density
VSEPR Theory
• Encountered geometries
– Electronic geometry – determined by the
location of “___” the regions of electron
density around the center atom(s)
– Molecular geometry – determined by the
arrangement of _____ only around the center
atom(s)
• The does not include lone electron pairs. The
molecular shape differs from the electron shape if
lone pairs are present
An example is H2O
VSEPR Theory
• Lone pairs of electrons occupy more space
than bonding pairs. As a consequence, there
is an order of the magnitude of repulsions
– lp/lp > lp/bp > bp/bp
As a result, the bond angles around a center
atoms can be distorted (reduced) from the
predicted values
CH4 and H2O What are the H-C-H and H-O-H
bond angles. If a change is observed, why?
Molecular Geometry and Polarity
• The polarity can be determine once the geometry
is known
• A polar bond is created if the atoms sharing the
electron pair have different electronegativities
– HCl and the associated dipole moment. This molecule
is polar. For diatomics, determination of polarity is
easy. What if the molecule has two or more atoms? All
the dipole have to be summed. If the sum equals zero,
the molecule has no dipole.
Molecular Geometry and Polarity
• A dipole moment (bond dipole) has _____ and
_____. Both must be considered when determines
if a molecule is polar.
– CO2 and H2O. Do these molecules have net dipoles?
• Conditions for polarity
– There must be at least one polar bond or lone pair on a
central atom
– The bond dipoles must not cancel or if there are two or
more lone pairs on the central atom, they must not be
arranged so that their polarities cancel
CO2, H2O, and O3
Molecular Geometry and Polarity
Valence Bond(VB) Theory
• VB theory describes how bonding occurs
• Describes how the atomic orbitals overlap to
produce the bonding geometry predicted by VSEPR
– Go back and review atomic orbitals if necessary
• Electrons are arranged in atomic orbitals according
to energy. The set of atomic orbitals, however, may
not be of lowest possible energy upon bonding
covalently to neighboring atoms.
Valence Bond(VB) Theory
• The valence shell orbitals (atomic orbitals)
commonly combine to change their character in
order to obtain a lower energy ‘mixed’ orbital set
for bonding in a particular geometry
– Which atomic orbitals would participate in bonding in
H, O, and C? These atomic orbitals can form a new set
of hybrid orbitals upon bonding.
• Hybrization – process by which ____ ____
combine to form a set of ‘mixed’ orbitals of lower
energy when bonding covalently
– The ‘mixed’ orbitals are called hybrid orbitals
Valence Bond(VB) Theory
• Hybrid orbitals on a center atom align themselves
with the bonding orbitals on the neighboring
atoms
– A ‘good overlap’ is necessary for sharing electrons in a
bond.
• Table 8-2 (refer to it)
– The label given to a set of hybridized orbitals reflects
the number and type of atomic orbitals used to produce
the set.
• Indicates the electronic geometry in agreement with VSEPR
Valence Bond(VB) Theory
Molecular Shapes and Bonding
• Simples structures will be analyzed based on
geometry type.
• Experimentally determined findings will be
discussed in light of these models.
• Terminology
– A – central atom
– B – atoms bonded to A
– U – lone pairs of electrons around A
AB3U represents three atoms bonded to a central atom
with one lone pair. An example would be NH3
Molecular Shapes and Bonding
Discussion sequence
• Experimental facts and Lewis formula
• VSEPR
– Electronic geometry
– Molecular geometry
– Polarity
• Valence bond theory
AB2 Molecules - No Lone Pairs
on A - Linear Molecules
• The BeCl2 molecule is linear and has melting
point of 405C.
– Draw BeCl2 and discuss electronic geometry
• Does the molecular geometry differ?
– The molecule does not satisfy the octet rule
– The compound bonds covalently due to the high charge
density on Be2+
• The electron cloud on the halide is distorted by the high charge
density
– BeBr2 and BeI2 also have linear geometries
AB2 Molecules - No Lone Pairs
on A - Linear Molecules
• The molecule possesses
two polar bonds (Be-Cl)
– EN = 1.5
• The molecule, however,
has no net dipole because
the two bond dipoles are
equal but in opposite
directions.
:Cl-Be-Cl :
Bond dipoles cancel.
This is a nonpolar
molecule.
AB2 Molecules - No Lone Pairs
on A - Linear Molecules
• Electronic Structures
1s 2s 2p
Be


3s
3p
Cl [Ne] 

Lewis Formulas
Be ··
·· .
·· Cl
··
The 2s orbital is full indicating that it will not bond. How will
the Be atom make these electrons available for bonding?
What happens in this molecule? Experimental data indicates
that the Be-Cl bonds are identical.
AB2 Molecules - No Lone Pairs
on A - Linear Molecules
• Valence Bond Theory (Hybridization)
1s 2s 2p 1s sp hyb
2p
Be 


3s
3p
Cl [Ne] 

The two atomic orbitals on Be hybridize to produce two sp
hybrid orbitals that have properties between the s and p
atomic orbitals. Notice that chlorine has a half-filled 3p
orbital that can overlap with the sp hybrid orbitals of Be.
AB2 Molecules - No Lone Pairs
on A - Linear Molecules
Two regions of electron density around the central atom
Illustrate how the sp orbitals overlap with the 3p orbitals on Cl
AB3 Molecules - No Lone Pairs
on A - Trigonal Planar Molecules
• Group IIIA elements that form covalent
compounds by bonding to three other atoms
– Octet rule is not satisfied but no big deal
• Boron trichloride is a trigonal molecule with a
melting point of -107C
– Does the molecular and electronic geometry
differ?
– The data indicates that this molecule is nonpolar
(no net dipole).
AB3 Molecules - No Lone Pairs
on A - Trigonal Planar Molecules
• Lewis structure predicts
trigonal planar geometry
• There are three bond dipoles
of equal length but different
direction.
– The bond dipoles cancel each
other
The molecule has no net dipole
How about BCl2H?
Cl
BCl
Cl
AB3 Molecules - No Lone Pairs
on A - Trigonal Planar Molecules
Lewis Formulas
.
B:
. Cl:
: :
• Electronic Structures
1s 2s 2p
B
 
3s
3p
Cl [Ne]   
Suppose that an electron in the 2s atomic orbital is promoted
to an empty 2p atomic orbital allowing for 3 unfilled atomic
orbitals for bonding. This would produce, however,
unequal energies for the three B-Cl bonds.
AB3 Molecules - No Lone Pairs
on A - Trigonal Planar Molecules
• Valence Bond Theory (Hybridization)
1s 2s 2p
1s sp2 hybrid
B  
 
3s
3p
Cl [Ne]   
The 2s and 2p atomic orbitals on B hybridize to
produce three sp orbitals (sp2 hybrid). Notice that
chlorine has a half-filled 3p orbital that can
overlap with the sp2 hybrid orbitals of Be.
AB3 Molecules - No Lone Pairs
on A - Trigonal Planar Molecules
Three regions of electron density around the central atom
Illustrate bonding with the Cl atoms on the hybridized B
AB4 Molecules - No Lone Pairs
on A - Tetrahedral Molecules
• Group IVA elements that form covalent
compounds by bonding to four other atoms
– Four electrons are shared and the octet rule is generally
satisfied
• CH4, methane, possesses a tetrahedral geometry
and has a melting point of -182C
– Would the molecular and electronic geometry differ?
– The data indicate that the molecule is nonpolar.
AB4 Molecules - No Lone Pairs
on A - Tetrahedral Molecules
• Lewis structure predicts
tetrahedral geometry
• There are four small bond
dipoles which cancel
– The molecule is nonpolar
• What about CCl3H and
CH3Cl?
– When the symmetry lowers, the
molecule becomes polar.
• Other molecules?
H
H
C
H
H
CH4
AB4 Molecules - No Lone Pairs
on A - Tetrahedral Molecules
• Electronic Structures
2s
2p
C [He]

1s
H
Lewis Formulas
.
C. :
H.
Suppose that an electron in the 2s atomic orbital is promoted
to an empty 2p atomic orbital allowing for 4 unfilled
atomic orbitals for bonding. This would produce,
however, unequal energies for the four C-H bonds.
AB4 Molecules - No Lone Pairs
on A - Tetrahedral Molecules
• Valence Bond
2s
2p
four sp3 hybrid orbitals
C [He] 
C [He]
1s
H
The 2s and 2p atomic orbitals on C hybridize to produce four
sp orbitals (sp3 hybrid). Notice that hyrogen has a halffilled 1s orbital that can overlap with the sp3 hybrid
orbitals of C.
Many AB4 type molecules have this hybridization
NH4+ is an AB4 type polyatomic ion
AB4 Molecules - No Lone Pairs
on A - Tetrahedral Molecules
Four regions of electron density around the central atom
Illustrate how the hydrogen atomic orbitals bond to the
hybridized carbon sp3 orbitals
Alkanes CnH2n+2
• alkanes are saturated hydrocarbons
• have the general formula CnH2n+2.
CH4 - methane
C2H6 or (H3C-CH3) - ethane
C3H8 or (H3C-CH2-CH3) - propane
• C atoms are located at the center of a tetrahedron
each alkane is a chain of interlocking tetrahedra
C atom at the center of each tetrahedron
enough H to form a total of four bonds for each C
AB3U Molecules - One Lone Pair
- Pyramidal Molecules
• Group VA elements (e.g. N) have five electrons in
the valence and commonly bond to three atoms
leaving a lone pair.
– The octet rule is satisfied
• The most common molecule is NH3.
– How many regions of electron density around nitrogen?
The bong angle is in this molecule is ~107. Why?
• Other common molecules are NF3, PF3, and the
polyatomic ion SO32-.
AB3U Molecules - One Lone Pair
- Pyramidal Molecules
• The Lewis structure predicts
tetrahedral electronic
geometry.
– Is the molecular geometry
different?
• There are three bond dipoles?
Detail.
– Is the molecule polar?
– How about NF3? How do the
polarities of the two molecules
compare (later)?
H
H
N
H
AB3U Molecules - One Lone Pair
- Pyramidal Molecules
Electronic Structures
2s
2p
N [He]

2s
2p
F [He]
  
1s
H
Lewis Formulas
There are three half-filled atomic orbitals on the nitrogen
(2p). The data suggests, however, that there are four nearly
equivalent orbitals (not three). Three orbitals are for
bonding and one for a lone pair.
AB3U Molecules - One Lone Pair
- Pyramidal Molecules
• Valence Bond
2s
N [He] 
2p

four sp3 hybrids
The 2s and 2p atomic orbitals hybridize to
form four sp3 hybrid orbitals. This
hybridization is also necessary to produce
the correct geometry for bonding.
Illustrate bonding with hydrogen 1s
atomic orbital.
Once again there are four regions of electron density around the
center atom
AB3U Molecules - One Lone Pair
- Pyramidal Molecules
• Let’s compare NH3 with NF3.
• The geometry of the both molecules is
already known
– Electronic geometry is _________
– Molecular geometry is _________
• How does the lone pair influence polarity?
It’s contribution has to be included to
determine polarity of a molecule.
AB3U Molecules - One Lone Pair
- Pyramidal Molecules
• The bond dipoles go opposite
directions on NH3 and NF3
– For NH3, the net dipole is enhanced
by the lone pair.
– For NF3, the net dipole is decreased
due to the lone pair
• Additionally, the H-N-H angle is
greater than the F-N-F angle due
to closer approach of the lone
pair to nitrogen on NF3
H
H
N
H
H-N-H = 107.3
N
F
F
F
F-N-F = 102.1
AB2U2 - Two Lone Pairs V-Shaped Molecules
• Group VIA elements (e.g. O) have six electrons in
the valence and commonly bond to two atoms
leaving two lone pairs.
– The octet rule is generally satisfied
• H2O is the most common molecule of this type.
– The molecular geometry is ______ and the electronic
geometry is _____
– Other examples of this type of molecule is H2S and
OCl2
AB2U2 - Two Lone Pairs V-Shaped Molecules
• The Lewis structure predicts
that the molecule is bent in
agreement with experimental
data.
– The actual H-O-H bond angle is
104.5 due to repulsions from
two lone pairs
• There are two bond dipoles
(O-H). Additionally, the net
dipole is enhanced by the lone
pairs.
– Illustrate
H
H
O
AB2U2 - Two Lone Pairs V-Shaped Molecules
Lewis Formulas
. O:
.
:
• Electronic Structures
2s
2p
O [He]


1s
H
H.
There are two half-filled atomic orbitals on the nitrogen (2p).
The data suggests, however, that there are four nearly
equivalent orbitals (not two). Two orbitals are for bonding
and two for lone pairs.
AB2U2 - Two Lone Pairs V-Shaped Molecules
• Valence Bond
2s
2p
four sp3 hybrids
O [He]  
 
The hybrid orbitals that are full belong to the lone pairs. The
half-filled orbitals are used for bonding.
Trigonal Bipyramidal Electronic
Geometry
• AB5, AB4U, AB3U2, and AB2U3
• Hybridization is sp3d.
• The lone pairs (if present) will arrange
themselves to minimize repulsive forces.
– lp/lp >> lp/bp > bp/bp
• This geometry is common for P, As, and Sb.
– All five valence electrons are shared (PF5)
Trigonal Bipyramidal Electronic
Geometry, AB5
• The VSEPR theory predicts
trigonal bipyramidal for the
electronic and molecular
geometry.
• There are three equatorial atoms
and to axial atoms. What are the
bond angles? Are the individual
bonds polar? Is the molecule
(type AB5) polar?
– Show molecule with this geometry.
··
·· F ·· ··
·
F
·
··
·· ··
F As ··
F ··
··
··
·· F ··
··
trigonal bipyramid
Trigonal Bipyramidal Electronic
Geometry, AB5
• Electronic Structures
3s 3p
P [Ne]

2s
2p
F [He]


Lewis Formulas
··
. P .
.
··
·· F
.
··
The 3d subshell is empty and participates in the rehybridization
(sp3d).
Trigonal Bipyramidal Electronic
Geometry, AB5
• Hybridization involves one d
orbital form the empty 3d
subshell and the 3s and 3d
orbitals.
– Illustrate from page 332
– Can also occur for n=4, 5, and 6
• There are no unshared pairs.
– 5 covalent bonds
This type of hybridization does
not occur for N. Why?
Trigonal Bipyramidal Electronic
Geometry with Lone Pairs
• AB4U, AB3U2, and
AB2U3
• Go through the
procedure for the
molecule, SF4.
– Where is the preferred
location of the lone pair?
• lp/lp>>lp/bp>bp/bp
– This molecular geometry
is termed as ______.
Is the molecule polar?
The Molecular Geometry, AB4U
Trigonal Bipyramidal Electronic
Geometry with Lone Pairs
• AB4U, AB3U2, and
AB2U3
• The ClF3 molecule
– Where are the likely
locations for the lone
pairs?
• The molecular geometry
is termed as _______.
What is the electronic
geometry?
The Molecular Geometry, AB3U2
Trigonal Bipyramidal Electronic
Geometry with Lone Pairs
• AB4U, AB3U2, and
AB2U3
• The I3- species
– Where are the likely
locations for the lone
pairs?
• The molecular
geometry is termed as
_______.
The Molecular Geometry, AB2U3.
Octahedral Electronic Geometry
• AB6, AB5U, and AB4U2
• Hybridization is sp3d2.
• Occurs for Group VIA elements below
oxygen.
• What are the predicted bond angles for this
geometry?
Octahedral Electronic Geometry:
AB6
• The VSEPR theory
predicts octahedral for the
electronic and molecular
geometry.
• Is this molecule polar?
• What are the bond
angles?
F
F
F
F
S
F
F
octahedral
Octahedral Electronic Geometry:
AB6
• Hybridization involves two d
orbital form the empty 3d
subshell and the 3s and 3p
orbitals.
– Illustrate from page 336
– Can also occur for n=4, 5, and 6
• There are no unshared pairs.
– 6 covalent bonds
Does this type of hybridization
occur for N?
Variations of Octahedral Shape
• If lone pairs are incorporated into the
octahedral structure, there are two possible
new shapes.
– One lone pair - square pyramidal
– Two lone pairs - square planar
• The resulting hybridization will be the
same.
Larger Molecules
• Cyclic molecules
• Linear and branched molecules
• Containing multiple types of elements