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CHAPTER 9 BondingMolecular Structure

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Title: CHAPTER 9 BondingMolecular Structure


1
CHAPTER 9 Bonding/Molecular Structure
Chemistry Chemical Reactivity Kotz Treichel
5th edition
Valence electrons versus core electrons Main
group Valence electrons are s and p electrons
in outermost shell. Core electrons are electrons
in inner shells including filled d subshells/ For
main group elements, the valence electrons equal
the group number.
2
CHAPTER 9 Bonding/Molecular Structure
Chemistry Chemical Reactivity Kotz Treichel
5th edition
Review of concepts in chapter Valence and core
electrons Lewis symbols for atoms Ionic and
covalent bonds Lewis structures and the octet
rule Resonance Formal charge Electronegativity
3
Valence Electrons (2)
For transition elements, valence electrons
include ns and (n-1)d orbitals. Table 9.1
Lewis electron dot symbols and octets Table 9.2
4
9.2 Chemical Bond Formation
Ionic bonds Lewis electron dot structure for
NaCl Figure 9.1
Covalent bonds Lewis electron dot structure for
Cl2
5
9.3 Bonding and Ionic Compounds
Figure 9.1 Exothermic reactions forming NaCl and
CaO from elements
Low ionization energies for metals halogens have
high affinity for electrons.
The structure of a compound, either ionic or
covalent is the one have the lowest potential
energy.
6
Ion Attraction and Energy
Formation of Na(g) is endothermic (496
kJ/mol) Formation of Cl-(g) is exothermic (-349
kJ/mol) Together, both processes are endothermic
(153 kJ/mol) Formation of the ion pair in the
ionic compound, however, is highly exothermic
(-498 kJ/mol). The energy associated with the
ionic attraction can be calculated using
Coulomb's law (remember Chapter 5). Dependent
upon magnitude of charges and distances
separating ions. (Figure 9.3) Lattice energies
and Table 9.3
7
9.4 Covalent Bonding and Lewis Structures
Covalent bonding and the use of lines to
illustrate shared electrons. (HH vs H-H)
Bond pair Lone pair (nonbonding electrons) Double
and triple bonds Octet rule (examples) Systematic
way to draw Lewis structures (example
CH2O) Example 9.2 Predicting Lewis
structures Table 9.4
8
The Lewis Structure of CO2
1) Decide on the central atom. Formulas are
often written in the order in which they are
connected (e.g. HOCl). Usually, the central atom
is less electronegative and is written first in
the formula (e.g. CCl4). Hydrogen is never the
central atom because in only forms one bond. For
many of the examples we will see, the central
atom will obvious because it will correspond to
the element that has only 1 atom in the formula.
O C O
9
The Lewis Structure of CO2 (2)
2) Sum the valence electrons from all atoms.
1 C atom 4 valence electrons 2 O atoms 2 x
6 valence electrons Total valence electrons 16
10
The Lewis Structure of CO2 (3)
3) Write the central atom and connect the other
atoms using bonds.
O C O
11
The Lewis Structure of CO2 (4a)
4a) Complete the octets for the atoms bonded to
the central atom.
12
The Lewis Structure of CO2 (4b)
4b) Place any leftover electrons on the central
atom. (for this structure 16 - 16 0 are
unused)
13
The Lewis Structure of CO2 (5)
5) If the central atom has less than 8
electrons, use the lone pairs on the
terminal atoms to form double or triple bonds.
14
The Lewis Structure of CCl4
1) Decide on the central atom. (C)
Cl Cl C Cl Cl
15
The Lewis Structure of CCl4
1) Decide on the central atom. (C)
2) Sum the valence electrons from all atoms.
1 C atom 4 valence electrons 4 Cl atoms 4
x 7 valence electrons Total valence electrons
32
16
The Lewis Structure of CCl4 (2)
3) Write the central atom and connect the other
atoms using bonds.
17
The Lewis Structure of CCl4 (3)
4a) Complete the octets for the atoms bonded to
the central atom.
18
The Lewis Structure of CCl4 (4)
4b) Place any leftover electrons on the central
atom. (for this structure 32 - 32 0 are
unused)
19
The Lewis Structure of H2O
1) Decide on the central atom. (O)
2) Sum the valence electrons from all atoms.
1 O atom 6 valence electrons 2 H atoms 2 x
1 valence electrons Total valence electrons 8
20
The Lewis Structure of H2O (2)
3) Write the central atom and connect the other
atoms using bonds.
H O H
21
The Lewis Structure of H2O (3)
4a) Complete the octets for the atoms bonded to
the central atom. Since hydrogen can accommodate
only two electrons, that step does not apply
here. (the previous structure remains unchanged)
H O H
22
The Lewis Structure of H2O (4)
4b) Place any leftover electrons on the central
atom. (for this structure 8 - 4 4 are unused)
23
9.4 Covalent Bonding and Lewis Structures(2)
Oxo acids and their anions (Table
9.5) Isoelectronic species (skip)
Resonance and resonance structures Ozone
(resonance hybrid) Benzene, carbonate ion,
hydrogen carbonate ion
24
The Lewis Structure of SF4
1) Decide on the central atom.
2) Sum the valence electrons from all atoms.
1 S atom 6 valence electrons 4 F atoms 4 x
7 valence electrons Total valence electrons 34
25
The Lewis Structure of SF4 (2)
3) Write the central atom and connect the other
atoms using bonds.
26
The Lewis Structure of SF4 (3)
4a) Complete the octets for the atoms bonded to
the central atom.
27
The Lewis Structure of SF4 (4)
4b) Place any leftover electrons on the central
atom. (for this structure 34 - 32 2 are
unused)
28
The Lewis Structure of SF4 (5)
Conclusions this molecule provides an example of
a structure that does not obey the octet rule.
Since sulfur occurs in the third row of the
periodic table, it is possible to form structures
with "expanded" valence shells.
Elements in the third row and beyond are able to
utilize empty d orbitals to accommodate
additional electrons and, therefore, form more
complex structures.
29
Other Exceptions to the Octet Rule
Compounds in which an atom has fewer than eight
valence electrons. e.g. boron trifluoride (BF3)
Molecules with an odd number of electrons. e.g.
NO2
30
Charge Distribution in Covalent Bonds
Formal charge group number - number of LPE
1/2(number BE) Bond polarity polar covalent
bond (dipolar Figure 9.7 (dipolar) Nonpolar
covalent bond Electronegativity Figure
9.9 Example 9.7 Combining formal charge and bond
polarity
31
Ambiguous Situations
For compounds such as boron trifluoride (BF3),
the octet rule may or may not be obeyed depending
upon the way the structure is drawn. The best
structure is often determined by using the formal
charge method. (see Example 9.8, page 352, to see
how this is applied to BF3).
Recall formal charge Group - number LPE
1/2 (number BE)
32
Lewis Structure of phosphate ion
What is the Lewis structure of phosphate ion
(PO43-)?
Or
33
Lewis Structure of phosphate ion (2)
Calculate formal charges.
For O 6 - 6 0.5(2) -1 For P 5 - 0
0.5(8) 1
34
Lewis Structure of phosphate ion (3)
Calculate formal charges.
For O 6 - 4 0.5(4) 0 For P 5 - 0
0.5(10) 0
35
Lewis Structure of phosphate ion (4)
What is the conclusion? One structure follows the
octet rule and the other minimizes the formal
charges. Depending upon the reference, you may
see both ways used. There are a number of other
chemical species that can be treated similarly
(e.g. SO42-, POCl3).
Note that oxygen usually forms 2 bonds in
molecules.
36
Molecular Shape Using VSEPR
The valence-shell electron-pair repulsion (VSEPR)
model provides a convenient way to identify the
three-dimensional shapes of molecules.
VSEPR looks at structures of the type AXn, where
is A is a central atom surrounded by electron
domains (or pairs (electron-rich regions).
37
Molecular Shape Using VSEPR (2)
An electron domain is one of the
following single bond nonbonding pair of
electrons multiple bond
38
Molecular Geometry with Balloons
Figure 9.11 Balloons are used to represent
electron pairs in the Valence Shell Electron Pair
Repulsion (VSEPR) model.
Balloons mimic electron domains in that they
repel one another and will attempt to separate as
much as possible. Based on the number of groups
around the central atom, several characteristic
geometries with respect to electron domains are
encountered.
39
Electron Pair Geometry vs Molecular Geometry
Although non-bonding electrons influence the
shape of molecules, only bonded atoms are used to
determine the molecular shape.
Figure 9.12 Geometries predicted for molecules
containing only single covalent bonds around the
central atom. Note the various bond angles.
Determine the molecular geometry of ammonia (NH3).
40
Molecular Geometry of Ammonia (NH3)
Use the following steps - determine the Lewis
structure - determine the electron-pair
geometry - determine the molecular geometry
Result ammonia is trigonal pyramidal.
41
Effect of Lone Pairs and Multiple Bonds
Lone pairs and multiple bonds tend to occupy more
space than single bonds (i.e. they acts like
bigger balloons). This leads to distortions of
the predicted bond angles relative to situations
when the electron pairs consist of single bonded
atoms.
Figure 9.13 Molecular geometries of methane,
ammonia and water.
42
Atoms with More Than Four Electron Pairs
With more than four electron pairs around a
central atom, the situation becomes more
complicated. For something like a
trigonal-bipyramidal geometry, the positions
around the central atom are not equivalent.
Treating the molecule as a sphere, we have
equatorial positions and axial positions. Lone
pairs, which require more space, prefer to occupy
the equatorial positions rather than the axial
positions (see figure)
43
9.10 Molecular Polarity
Recall that when two atoms with different
electronegativities are covalently bonded, they
form a polar bond. There is an imbalance of
charge that depends on the electronegativity
difference.
H Cl ?EN 3.0 - 2.1 0.9 Cl Cl ?EN 3.0
- 3.0 0 (nonpolar)
44
9.10 Molecular Polarity (2)
The term "polar" implies a bond that has a
partial positive end and a partial negative end.
Molecules that have polar bonds can also be
considered polar as a whole. For a linear
molecule such as HCl, one side of the molecule is
positive and the other side is negative. (see
Figure 9.15a)
Figure 9.15b Polar molecules in an electric
field.
45
Dipole Moments
The dipole moment is a measure of the polarity of
a molecule.
Dipole moment (?) product of the magnitude of
the partial charges (? and -?) and the distance
by which they are separated.
Units of dipole moment coulomb-meter or debye
(D) 1D 3.34 x 10-30 Cm
46
Dipole Moments (2)
Table 9.10 Dipole moments for selected molecules.
Note that as the dipole moment increases, the
polarity of the molecule also increases. A
dipole moment of zero means that the molecule is
nonpolar. From a practical standpoint, if the
dipole moment is close to zero, the molecule is
also considered to be nonpolar.
47
Determining Molecular Polarity
To determine if a molecule is polar, you first
have to determine if it has polar bonds. If
polar bonds are present, polarity of the molecule
will depend upon the molecular geometry and the
symmetry of the molecule.
Polar bonds on opposite ends of a molecule can
offset the effects of each other and lead to a
situation in which there is no net dipole moment.
48
Determining Molecular Polarity (2)
Essentially, polar bonds act as vectors. Vectors
have both magnitude and direction and can be
added together geometrically. Depending upon the
orientation and magnitude, they can sometimes
cancel each other out so that there is no net
imbalance of charge for the molecule.
The analogy of people pulling on a rope is
sometimes used. Other visualization techniques
will be shown in class.
49
Examples of Molecular Polarity
Figure 9.16 CO2 and H2O find centers of
positive and negative charge Figure 9.17 BF3,
Cl2CO and NH3 Example 9.14 NF3, CH2Cl2 and
SF4 Compare CH4 and CHCl3
Cooking with Microwaves (page 367). Note
microwave spectroscopy is one way to determine
dipole moments experimentally.
50
CHAPTER 10 Orbital Hybridization
Chemistry Chemical Reactivity Kotz Treichel
5th edition
Review of concepts in chapter
51
Chapter 10 Orbital Hybridization
Covalent bonds can be understood by looking at
the way in which atomic orbitals from adjacent
atoms interact. Valence bond theory uses quantum
mechanics to help explain why bonds exist between
atoms..
Provides a model for explaining the equivalency
of bonds that occur within a molecule and can be
used in conjunction with VSEPR to explain
geometries.
52
Valence Bond Theory
Bonds form between atoms when a valence atomic
orbital of one atom merges with that of another.
The orbitals share a region of space or overlap.
The overlap of orbitals allows two electrons of
opposite spin to share a common space and to form
a covalent bond.
Figure 10.1 shows what happens when 2 hydrogen
atoms are brought together to form a bond.
53
Figure 10.1 (continued)
Figure 10.1 shows the a graph of potential energy
versus internuclear distance (between 2 H atoms).
When the H atoms are brought closer together, the
electron from one atom is attracted to the
nucleus of the other atom.
The attractive forces distort the electron clouds
associated with the electrons and the potential
energy of the system is lowered.
For H, the potential energy is at a minimum when
the distance between nuclei is 74 pm.
54
Figure 10.1 (2)
When the electrons from each H atom pair up to
make a bond, there is a net stabilization (lower
potential energy) that results. This energy
difference can be calculated and the value
approximates the experimentally determined bond
energy.
The two atoms experience orbital overlap, there
is an increased probability of finding the
bonding electrons in the region of space between
the nuclei.
55
Orbital Overlap and Sigma Bonds
When the two H nuclei move closer together,
repulsion effects increase the potential energy
and the molecule becomes less stable.
The type of bond formed in H2 is called a sigma
(?) bond. The electron density for a ? bond is
greatest along the axis of the bond.
See summary of valence bond approach on page 383.
Figure 10.2 Covalent bond formation in H2, HF,
and F2.
56
Hybridization
The simple orbital overlap concept leads to
difficulties for molecules involving more atoms.
The need to have more unpaired electrons for
bonding and the desire to have bonding
arrangements that are equivalent leads to the
idea of hybridization.
Hybridization involves the mixing of atomic
orbitals on one atom to form new orbitals called
hybrid orbitals.
57
Characteristics of Hybrid Orbitals
Hybrid orbitals result from the mixing of s, p
and d orbitals.
The number of hybrid orbitals is always the same
as the number of atomic orbitals used to create
the hybrid set.
Hybrid orbitals are directed more towards the
terminal atoms than unhybridized orbitals. This
leads to better orbital overlap and stronger
bonds.
58
Sets of Hybrid Orbitals
The number of hybrid orbitals required by an atom
in a molecule or ion depends upon the
electron-pair geometry. A hybrid orbital is
required for each electron pair on the central
atom (lone pair or bond).
Figure 10.5 Hybrid orbitals for two through six
electron pairs.
59
Valence Bond Theory Examples
Examples for tetrahedral geometries Figure 10.6
Methane (CH4) Figure 10.7 Ammonia (NH3) Figure
10.8 Water (H2O)
60
Valence Bond Theory Examples (2)
Examples for trigonal planar geometries BF3,
CO32-, H2CCH2 Requires 3 hybrid orbitals in a
plane (120? apart) Need to combine 1 s and 2 p
orbitals.
For BF3, each B-F bond results from the overlap
of an sp2 orbital on boron with a p orbital on
the fluorine. One p orbital is unused and is not
occupied by electrons. (see geometry in Figure
10.9)
61
Hybridization In BF3
??
?
2s 2p
?
?
?
2s 2p
?
?
?
sp2 2p
62
Valence Bond Theory Examples (3)
Examples for linear geometries BeCl2 sp
hybridization with 2 p orbitals unoccupied linear
geometry see Figure 10.10 for geometry
63
Valence Bond Theory Examples (4)
Examples for requiring the use of d orbitals PF5,
SF6 Example 10.3 PF5 - geometry is trigonal
bipyramidal - hybridization is sp3d
64
Hybridization Involving d orbitals
65
Multiple Bonds
Multiple bonds involve the overlap of more than
one electron pair. Double bonds require two sets
of overlapping orbitals and two electron pairs.
Triple bonds require three.
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