Taste The taste of a food depends on the interaction between
the food molecules and taste cells on your tongue The main factors
that affect this interaction are the shape of the molecule and
charge distribution within the molecule The food molecule must fit
snugly into the active site of specialized proteins on the surface
of taste cells When this happens, changes in the protein structure
cause a nerve signal to transmit
Slide 3
Sugar & Artificial Sweeteners Sugar molecules fit into the
active site of taste cell receptors called Tlr3 receptor proteins
When the sugar molecule (the key) enters the active site (the
lock), the different subunits of the T1r3 protein split apart This
split causes ion channels in the cell membrane to open, resulting
in nerve signal transmission Artificial sweeteners also fit into
the Tlr3 receptor, sometimes binding to it even stronger than sugar
making them sweeter than sugar
Slide 4
Structure Determines Properties! Properties of molecular
substances depend on the structure of the molecule The structure
includes many factors, such as: the skeletal arrangement of the
atoms the kind of bonding between the atoms ionic, polar covalent,
or covalent the shape of the molecule Bonding theory should allow
you to predict the shapes of molecules
Slide 5
Molecular Geometry Molecules are 3-dimensional objects We often
describe the shape of a molecule with terms that relate to
geometric figures These geometric figures have characteristic
corners that indicate the positions of the surrounding atoms around
a central atom in the center of the geometric figure The geometric
figures also have characteristic angles that we call bond
angles
Slide 6
Lewis Theory Predicts Electron Groups Lewis theory predicts
there are regions of electrons in an atom Some regions result from
placing shared pairs of valence electrons between bonding nuclei
Other regions result from placing unshared valence electrons on a
single nuclei
Slide 7
Using Lewis Theory to Predict Molecular Shapes Lewis theory
says that these regions of electron groups should repel each other
because they are regions of negative charge This idea can then be
extended to predict the shapes of molecules the position of atoms
surrounding a central atom will be determined by where the bonding
electron groups are the positions of the electron groups will be
determined by trying to minimize repulsions between them
Slide 8
VSEPR Theory Electron groups around the central atom will be
most stable when they are as far apart as possible we call this
valence shell electron pair repulsion theory because electrons are
negatively charged, they should be most stable when they are
separated as much as possible The resulting geometric arrangement
will allow us to predict the shapes and bond angles in the
molecule
Slide 9
Slide 10
Electron Groups The Lewis structure predicts the number of
valence electron pairs around the central atom(s) Each lone pair of
electrons constitutes one electron group on a central atom Each
bond constitutes one electron group on a central atom regardless of
whether it is single, double, or triple ONO there are three
electron groups on N three lone pair one single bond one double
bond
Slide 11
Electron Group Geometry There are five basic arrangements of
electron groups around a central atom based on a maximum of six
bonding electron groups though there may be more than six on very
large atoms, it is very rare Each of these five basic arrangements
results in five different basic electron geometries in order for
the molecular shape and bond angles to be a perfect geometric
figure, all the electron groups must be bonds and all the bonds
must be equivalent For molecules that exhibit resonance, it doesnt
matter which resonance form you use the electron geometry will be
the same Tro: Chemistry: A Molecular Approach, 2/e
Slide 12
Linear When there are two electron groups around the central
atom, they will occupy positions on opposite sides of the central
atom This results in the electron groups taking a linear geometry
The bond angle is 180
Slide 13
Trigonal Planar When there are three electron groups around the
central atom, they will occupy positions in the shape of a triangle
around the central atom This results in the electron groups taking
a trigonal planar geometry The bond angle is 120
Slide 14
Tetrahedral When there are four electron groups around the
central atom, they will occupy positions in the shape of a
tetrahedron around the central atom This results in the electron
groups taking a tetrahedral geometry The bond angle is 109.5
Slide 15
Trigonal Bipyramidal 5 electron groups around a central atom
occupy positions in the shape of two tetrahedra base-to-base with
the central atom in the center of the shared bases The positions
above and below the central atom are called the axial positions The
positions in the same base plane as the central atom are called the
equatorial positions The bond angle between equatorial positions is
120 The bond angle between axial and equatorial positions is
90
Slide 16
Octahedral When there are six electron groups around the
central atom, they will occupy positions in the shape of two
square-base pyramids that are base-to-base with the central atom in
the center of the shared bases This results in the electron groups
taking an octahedral geometry it is called octahedral because the
geometric figure has eight sides All positions are equivalent The
bond angle is 90
Slide 17
Molecular Geometry The actual geometry of the molecule may be
different from the electron geometry When the electron groups are
attached to atoms of different size, or when the bonding to one
atom is different than the bonding to another, this will affect the
molecular geometry around the central atom Lone pairs also affect
the molecular geometry they occupy space on the central atom, but
are not seen as points on the molecular geometry
Slide 18
Not Quite Perfect Geometry Because the bonds and atom sizes are
not identical in formaldehyde, the observed angles are slightly
different from ideal
Slide 19
The Effect of Lone Pairs Lone pair groups occupy more space on
the central atom because their electron density is exclusively on
the central atom rather than shared like bonding electron groups
Relative sizes of repulsive force interactions is Lone Pair Lone
Pair > Lone Pair Bonding Pair > Bonding Pair Bonding Pair
This affects the bond angles, making the bonding pair bonding pair
angles smaller than expected
Slide 20
Effect of Lone Pairs The bonding electrons are shared by two
atoms, so some of the negative charge is removed from the central
atom The nonbonding electrons are localized on the central atom, so
area of negative charge takes more space
Slide 21
Bond Angle Distortion from Lone Pairs
Slide 22
Slide 23
Bent Molecular Geometry: Derivative of Trigonal Planar Electron
Geometry When there are three electron groups around the central
atom, and one of them is a lone pair, the resulting shape of the
molecule is called a trigonal planar bent shape The bond angle is
less than 120 because the lone pair takes up more space
Slide 24
Pyramidal Geometry: A Derivative of Tetrahedral Electron
Geometry When there are four electron groups around the central
atom, and one is a lone pair, the result is called a pyramidal
shape because it is a triangular-base pyramid with the central atom
at the apex Bond angle is less than 109.5
Slide 25
Bent Tetrahedral Molecular Geometry: A Derivative of
Tetrahedral Electron Geometry When there are four electron groups
around the central atom, and two are lone pairs, the result is
called a tetrahedral bent shape it is planar it looks similar to
the trigonal planarbent shape, except the angles are smaller the
bond angle is less than 109.5 Tro: Chemistry: A Molecular Approach,
2/e
Slide 26
Derivatives of the Trigonal Bipyramidal Electron Geometry When
there are five electron groups around the central atom, and some
are lone pairs, they will occupy the equatorial positions because
there is more room Seesaw Shape - five electron groups around the
central atom, and one is a lone pair ( aka distorted tetrahedron)
T-shaped 5 electron groups around the central atom, and two are
lone pairs Linear shape 5 electron groups around the central atom,
and three are lone pairs The bond angles between equatorial
positions are less than 120 The bond angles between axial and
equatorial positions are less than 90 linear = 180
axialtoaxial
Slide 27
Replacing Atoms with Lone Pairs in the Trigonal Bipyramid
System
Slide 28
Seesaw Shape
Slide 29
TShape
Slide 30
Slide 31
Linear Shape
Slide 32
Derivatives of the Octahedral Geometry When there are six
electron groups around the central atom, and some are lone pairs,
each even number lone pair will take a position opposite the
previous lone pair Square Pyramid Shape 6 electron groups around
the central atom, and one is a lone pair, the result is called a
the bond angles between axial and equatorial positions is less than
90 Square Planar Shape 6 electron groups around the central atom,
and two are lone pairs, the result is called a the bond angles
between equatorial positions is 90
Slide 33
Square Pyramidal Shape
Slide 34
Square Planar Shape
Slide 35
Slide 36
Predicting the Shapes Around Central Atoms 1.Draw the Lewis
structure 2.Determine the number of electron groups around the
central atom 3.Classify each electron group as bonding or lone
pair, and count each type remember, multiple bonds count as one
group 4. Use Table 10.1 to determine the shape and bond angles
Slide 37
Example 10.2: Predict the geometry and bond angles of PCl 3
1.Draw the Lewis structure a)26 valence electrons 2. Determine the
Number of electron groups around central atom a)four electron
groups around P
Slide 38
Example 10.2: Predict the geometry and bond angles of PCl 3
3.Classify the electron groups a)three bonding groups b)one lone
pair 4.Use Table 10.1 to determine the shape and bond angles a)four
electron groups around P = tetrahedral electron geometry b)three
bonding + one lone pair = trigonal pyramidal molecular geometry
c)trigonal pyramidal = bond angles less than 109.5
Slide 39
Practice Predict the molecular geometry and bond angles in SiF
5
Slide 40
Practice Predict the molecular geometry and bond angles in SiF
5 Si = 4e F 5 = 5(7e ) = 35e () = 1e total = 40e 5 electron groups
on Si 5 bonding groups 0 lone pairs Shape = trigonal bipyramid Bond
angles F eq SiF eq = 120 F eq SiF ax = 90 Si least electronegative
Si is central atom
Slide 41
Practice Predict the molecular geometry and bond angles in ClO
2 F
Slide 42
Cl = 7e O 2 = 2(6e ) = 12e F = 7e Total = 26e 4 electron groups
on Cl 3 bonding groups 1 lone pair Shape = trigonal pyramidal Bond
angles OClO < 109.5 OClF < 109.5 Cl least electronegative Cl
is central atom
Slide 43
Representing 3-Dimensional Shapes on a 2-Dimensional Surface
One of the problems with drawing molecules is trying to show their
dimensionality By convention, the central atom is put in the plane
of the paper Put as many other atoms as possible in the same plane
and indicate with a straight line For atoms in front of the plane,
use a solid wedge For atoms behind the plane, use a hashed
wedge
Slide 44
Slide 45
SF 6 S F F F FF F
Slide 46
Multiple Central Atoms Many molecules have larger structures
with many interior atoms We can think of them as having multiple
central atoms When this occurs, we describe the shape around each
central atom in sequence shape around left C is tetrahedral shape
around center C is trigonal planar shape around right O is
tetrahedral-bent
Slide 47
Describing the Geometry of Methanol
Slide 48
Describing the Geometry of Glycine
Slide 49
Practice Predict the molecular geometries in H 3 BO 3
Slide 50
3 electron groups on B B has 3 bonding groups 0 pone pairs 4
electron groups on O O has 2 bonding groups 2 lone pairs Practice
Predict the molecular geometries in H 3 BO 3 B = 3e O 3 = 3(6e ) =
18e H 3 = 3(1e ) = 3e Total = 24e Shape on B = trigonal planar B
least electronegative B Is Central Atom oxyacid, so H attached to O
Shape on O = tetrahedral bent
Slide 51
Polarity of Molecules For a molecule to be polar it must 1.have
polar bonds electronegativity difference - theory bond dipole
moments - measured 2.have an unsymmetrical shape vector addition
Polarity affects the intermolecular forces of attraction therefore
boiling points and solubilities like dissolves like Nonbonding
pairs affect molecular polarity, strong pull in its direction
Slide 52
Molecule Polarity The HCl bond is polar. The bonding electrons
are pulled toward the Cl end of the molecule. The net result is a
polar molecule.
Slide 53
Vector Addition
Slide 54
Slide 55
Molecule Polarity The OC bond is polar. The bonding electrons
are pulled equally toward both O ends of the molecule. The net
result is a nonpolar molecule.
Slide 56
Molecule Polarity The HO bond is polar. Both sets of bonding
electrons are pulled toward the O end of the molecule. The net
result is a polar molecule.
Slide 57
Predicting Polarity of Molecules 1.Draw the Lewis structure and
determine the molecular geometry 2.Determine whether the bonds in
the molecule are polar a)if there are not polar bonds, the molecule
is nonpolar 3.Determine whether the polar bonds add together to
give a net dipole moment
Slide 58
Example 10.5: Predict whether NH 3 is a polar molecule 1.Draw
the Lewis structure and determine the molecular geometry a)eight
valence electrons b)three bonding + one lone pair = trigonal
pyramidal molecular geometry
Slide 59
Example 10.5: Predict whether NH 3 is a polar molecule
2.Determine if the bonds are polar a)electronegativity difference
b)if the bonds are not polar, we can stop here and declare the
molecule will be nonpolar EN N = 3.0 EN H = 2.1 3.0 2.1 = 0.9
therefore the bonds are polar covalent
Slide 60
Example 10.5: Predict whether NH 3 is a polar molecule
3)Determine whether the polar bonds add together to give a net
dipole moment a)vector addition b)generally, asymmetric shapes
result in uncompensated polarities and a net dipole moment The HN
bond is polar. All the sets of bonding electrons are pulled toward
the N end of the molecule. The net result is a polar molecule.
Slide 61
Practice Decide whether the following molecules are polar EN O
= 3.5 N = 3.0 Cl = 3.0 S = 2.5
Slide 62
Practice Decide whether the following molecules Are polar polar
nonpolar 1. polar bonds, N-O 2. asymmetrical shape 1. polar bonds,
all S-O 2. symmetrical shape Trigonal Bent Trigonal Planar 2.5
Slide 63
Molecular Polarity Affects Solubility in Water Polar molecules
are attracted to other polar molecules Because water is a polar
molecule, other polar molecules dissolve well in water and ionic
compounds as well Some molecules have both polar and nonpolar
parts
Slide 64
Problems with Lewis Theory Lewis theory generally predicts
trends in properties, but does not give good numerical predictions
e.g. bond strength and bond length Lewis theory gives good first
approximations of the bond angles in molecules, but usually cannot
be used to get the actual angle Lewis theory cannot write one
correct structure for many molecules where resonance is important
Lewis theory often does not predict the correct magnetic behavior
of molecules e.g. O 2 is paramagnetic, though the Lewis structure
predicts it is diamagnetic
Slide 65
Valence Bond Theory Linus Pauling and others applied the
principles of quantum mechanics to molecules They reasoned that
bonds between atoms would occur when the orbitals on those atoms
interacted to make a bond The kind of interaction depends on
whether the orbitals align along the axis between the nuclei, or
outside the axis
Slide 66
Orbital Interaction As two atoms approached, the half-filled
valence atomic orbitals on each atom would interact to form
molecular orbitals molecular orbtials are regions of high
probability of finding the shared electrons in the molecule The
molecular orbitals would be more stable than the separate atomic
orbitals because they would contain paired electrons shared by both
atoms the potential energy is lowered when the molecular orbitals
contain a total of two paired electrons compared to separate one
electron atomic orbitals
Slide 67
Orbital Diagram for the Formation of H 2 S + H 1s1s HS bond H
1s1s S 3s3s3p3p HS bond Predicts bond angle = 90 Actual bond angle
= 92
Slide 68
Valence Bond Theory Hybridization One of the issues that arises
is that the number of partially filled or empty atomic orbitals did
not predict the number of bonds or orientation of bonds C = 2s 2 2p
x 1 2p y 1 2p z 0 would predict two or three bonds that are 90
apart, rather than four bonds that are 109.5 apart To adjust for
these inconsistencies, it was postulated that the valence atomic
orbitals could hybridize before bonding took place one
hybridization of C is to mix all the 2s and 2p orbitals to get four
orbitals that point at the corners of a tetrahedron
Slide 69
Unhybridized C Orbitals Predict the Wrong Bonding &
Geometry
Slide 70
Valence Bond Theory Main Concepts 1.The valence electrons of
the atoms in a molecule reside in quantum-mechanical atomic
orbitals. The orbitals can be the standard s, p, d, and f orbitals,
or they may be hybrid combinations of these. 2.A chemical bond
results when these atomic orbitals interact and there is a total of
two electrons in the new molecular orbital a)the electrons must be
spin paired 3.The shape of the molecule is determined by the
geometry of the interacting orbitals
Slide 71
Hybridization Some atoms hybridize their orbitals to maximize
bonding more bonds = more full orbitals = more stability
Hybridizing is mixing different types of orbitals in the valence
shell to make a new set of degenerate orbitals sp, sp 2, sp 3, sp 3
d, sp 3 d 2 Same type of atom can have different types of
hybridization C = sp, sp 2, sp 3
Slide 72
Hybrid Orbitals The number of standard atomic orbitals combined
= the number of hybrid orbitals formed combining a 2s with a 2p
gives two 2sp hybrid orbitals H cannot hybridize!! its valence
shell only has one orbital The number and type of standard atomic
orbitals combined determines the shape of the hybrid orbitals The
particular kind of hybridization that occurs is the one that yields
the lowest overall energy for the molecule
sp 3 Hybridization Atom with four electron groups around it
tetrahedral geometry 109.5 angles between hybrid orbitals Atom uses
hybrid orbitals for all bonds and lone pairs
Slide 75
Slide 76
Orbital Diagram of the sp 3 Hybridization of C
Slide 77
sp 3 Hybridized Atoms Orbital Diagrams Unhybridized atom 2s2s
2p2p sp 3 hybridized atom 2sp 3 C 2s2s 2p2p 2sp 3 N Place electrons
into hybrid and unhybridized valence orbitals as if all the
orbitals have equal energy Lone pairs generally occupy hybrid
orbitals
Slide 78
Practice Draw the orbital diagram for the sp 3 hybridization of
each atom Unhybridized atom 3s3s 3p3p Cl 2s2s 2p2p O sp 3
hybridized atom 3sp 3 2sp 3
Slide 79
Bonding with Valence Bond Theory According to valence bond
theory, bonding takes place between atoms when their atomic or
hybrid orbitals interact overlap To interact, the orbitals must
either be aligned along the axis between the atoms, or The orbitals
must be parallel to each other and perpendicular to the interatomic
axis
Slide 80
Methane Formation with sp 3 C
Slide 81
Ammonia Formation with sp 3 N
Slide 82
Types of Bonds A sigma ( ) bond results when the interacting
atomic orbitals point along the axis connecting the two bonding
nuclei either standard atomic orbitals or hybrids stos, ptop,
hybridtohybrid, stohybrid, etc. A pi ( ) bond results when the
bonding atomic orbitals are parallel to each other and
perpendicular to the axis connecting the two bonding nuclei between
unhybridized parallel p orbitals The interaction between parallel
orbitals is not as strong as between orbitals that point at each
other; therefore bonds are stronger than bonds
Slide 83
Slide 84
Orbital Diagrams of Bonding Overlap between a hybrid orbital on
one atom with a hybrid or nonhybridized orbital on another atom
results in a bond Overlap between unhybridized p orbitals on bonded
atoms results in a bond
Slide 85
CH 3 NH 2 Orbital Diagram sp 3 C sp 3 N 1s H
Slide 86
Formaldehyde, CH 2 O Orbital Diagram sp 2 C sp 2 O 1s H p Cp
O
Slide 87
sp 2 Atom with three electron groups around it trigonal planar
system C = trigonal planar N = trigonal bent O = linear 120 bond
angles flat Atom uses hybrid orbitals for bonds and lone pairs,
uses nonhybridized p orbital for bond
Slide 88
Slide 89
sp 2 Hybridized Atoms Orbital Diagrams Unhybridized atom 2s2s
2p2p sp 2 hybridized atom 2sp 2 2p2p C 3 1 2s2s 2p2p 2sp 2 2p2p N 2
1
Slide 90
Practice Draw the orbital diagram for the sp 2 hybridization of
each atom. How many and bonds would you expect each to form?
Unhybridized atom 2s2s 2p2p B 3 0 2s2s 2p2p O 1 1 sp 2 hybridized
atom 2sp 2 2p2p 2p2p
Slide 91
Hybrid orbitals overlap to form a bond. Unhybridized p orbitals
overlap to form a bond.
Slide 92
CH 2 NH Orbital Diagram sp 2 C sp 2 N 1s H p Cp N CN H HH
Slide 93
Bond Rotation Orbitals form the bond point along the
internuclear axis rotation around that bond does not require
breaking the interaction between the orbitals Orbitals that form
the bond interact above and below the internuclear axis rotation
around the axis requires the breaking of the interaction between
the orbitals
Slide 94
Slide 95
Slide 96
sp Atom with two electron groups linear shape 180 bond angle
Atom uses hybrid orbitals for bonds or lone pairs, uses
nonhybridized p orbitals for bonds
Slide 97
Slide 98
Slide 99
sp Hybridized Atoms Orbital Diagrams Unhybridized atom 2s2s
2p2p sp hybridized atom 2sp 2p2p C22C22 2s2s 2p2p 2sp 2p2p
N12N12
Slide 100
HCN Orbital Diagram sp C sp N 1s H s p Cp N 22
Slide 101
sp 3 d Atom with five electron groups around it trigonal
bipyramid electron geometry Seesaw, TShape, Linear 120 & 90
bond angles Use empty d orbitals from valence shell d orbitals can
be used to make bonds
Slide 102
Slide 103
sp 3 d Hybridized Atoms Orbital Diagrams Unhybridized atom 3s3s
3p3p sp 3 d hybridized atom 3sp 3 d S 3s3s 3p3p 3sp 3 d P 3d3d 3d3d
(non-hybridizing d orbitals not shown)
Slide 104
SOF 4 Orbital Diagram sp 3 d S sp 2 O 2p F d Sp O 2p F
Slide 105
sp 3 d 2 Atom with six electron groups around it octahedral
electron geometry Square Pyramid, Square Planar 90 bond angles Use
empty d orbitals from valence shell to form hybrid d orbitals can
be used to make bonds
Slide 106
Slide 107
sp 3 d 2 Hybridized Atoms Orbital Diagrams Unhybridized atom sp
3 d 2 hybridized atom S 3s3s 3p3p 3sp 3 d 2 3d3d I 5s5s 5p5p 5d5d
5sp 3 d 2 (non-hybridizing d orbitals not shown)
Slide 108
Slide 109
Predicting Hybridization and Bonding Scheme 1.Start by drawing
the Lewis structure 2.Use VSEPR Theory to predict the electron
group geometry around each central atom 3.Use Table 10.3 to select
the hybridization scheme that matches the electron group geometry
4.Sketch the atomic and hybrid orbitals on the atoms in the
molecule, showing overlap of the appropriate orbitals 5.Label the
bonds as or
Slide 110
Example 10.7: Predict the hybridization and bonding scheme for
CH 3 CHO Draw the Lewis structure Predict the electron group
geometry around inside atoms C1 = 4 electron areas C1= tetrahedral
C2 = 3 electron areas C2 = trigonal planar
Slide 111
Example 10.7: Predict the hybridization and bonding scheme for
CH 3 CHO Determine the hybridization of the interior atoms C1 =
tetrahedral C1 = sp 3 C2 = trigonal planar C2 = sp 2 Sketch the
molecule and orbitals
Slide 112
Label the bonds Example 10.7: Predict the hybridization and
bonding scheme for CH 3 CHO
Slide 113
Practice Predict the hybridization of all the atoms in H 3 BO 3
H = cant hybridize B = 3 electron groups = sp 2 O = 4 electron
groups = sp 3
Slide 114
Practice Predict the hybridization and bonding scheme of all
the atoms in NClO N = 3 electron groups = sp 2 O = 3 electron
groups = sp 2 Cl = 4 electron groups = sp 3 ONCl Nsp 2 Clp Osp 2
Nsp 2 OpNp
Slide 115
Problems with Valence Bond Theory VB theory predicts many
properties better than Lewis theory bonding schemes, bond
strengths, bond lengths, bond rigidity However, there are still
many properties of molecules it doesnt predict perfectly magnetic
behavior of O 2 In addition, VB theory presumes the electrons are
localized in orbitals on the atoms in the molecule it doesnt
account for delocalization
Slide 116
Molecular Orbital Theory In MO theory, we apply Schrdingers
wave equation to the molecule to calculate a set of molecular
orbitals in practice, the equation solution is estimated we start
with good guesses from our experience as to what the orbital should
look like then test and tweak the estimate until the energy of the
orbital is minimized In this treatment, the electrons belong to the
whole molecule so the orbitals belong to the whole molecule
delocalization
Slide 117
LCAO The simplest guess starts with the atomic orbitals of the
atoms adding together to make molecular orbitals this is called the
Linear Combination of Atomic Orbitals method weighted sum Because
the orbitals are wave functions, the waves can combine either
constructively or destructively
Slide 118
Molecular Orbitals When the wave functions combine
constructively, the resulting molecular orbital has less energy
than the original atomic orbitals it is called a Bonding Molecular
Orbital , most of the electron density between the nuclei When the
wave functions combine destructively, the resulting molecular
orbital has more energy than the original atomic orbitals it is
called an Antibonding Molecular Orbital *, * most of the electron
density outside the nuclei nodes between nuclei
Slide 119
Interaction of 1s Orbitals
Slide 120
Molecular Orbital Theory Electrons in bonding MOs are
stabilizing lower energy than the atomic orbitals Electrons in
antibonding MOs are destabilizing higher in energy than atomic
orbitals electron density located outside the internuclear axis
electrons in antibonding orbitals cancel stability gained by
electrons in bonding orbitals
Slide 121
Energy Comparisons of Atomic Orbitals to Molecular
Orbitals
Slide 122
MO and Properties Bond Order = difference between number of
electrons in bonding and antibonding orbitals only need to consider
valence electrons may be a fraction higher bond order = stronger
and shorter bonds if bond order = 0, then bond is unstable compared
to individual atoms and no bond will form A substance will be
paramagnetic if its MO diagram has unpaired electrons if all
electrons paired it is diamagnetic
Slide 123
1s1s1s1s ** Hydrogen Atomic Orbital Hydrogen Atomic Orbital
Dihydrogen, H 2 Molecular Orbitals Because more electrons are in
bonding orbitals than are in antibonding orbitals, net bonding
interaction
Slide 124
H2H2 * Antibonding MO LUMO bonding MO HOMO
Slide 125
1s1s1s1s ** Helium Atomic Orbital Helium Atomic Orbital
Dihelium, He 2 Molecular Orbitals Because there are as many
electrons in antibonding orbitals as in bonding orbitals, there is
no net bonding interaction BO = (2-2) = 0
Slide 126
1s1s1s1s Lithium Atomic Orbitals Lithium Atomic Orbitals
Dilithium, Li 2 Molecular Orbitals Because more electrons are in
bonding orbitals than are in antibonding orbitals, there is a net
bonding interaction 2s2s2s2s Any fill energy level will generate
filled bonding and antibonding MOs; therefore only need to consider
valence shell BO = (4-2) = 1
Slide 127
bonding MO HOMO * Antibonding MO LUMO Li 2
Slide 128
Interaction of p Orbitals
Slide 129
Slide 130
Slide 131
O2O2 Dioxygen is paramagnetic Paramagnetic material has
unpaired electrons Neither Lewis theory nor valence bond theory
predict this result
Slide 132
O 2 as Described by Lewis and VB Theory
Slide 133
Oxygen Atomic Orbitals Oxygen Atomic Orbitals 2s2s2s2s 2p2p2p2p
Because more electrons are in bonding orbitals than are in
antibonding orbitals, there is a net bonding interaction Because
there are unpaired electrons in the antibonding orbitals, O 2 is
predicted to be paramagnetic O 2 MOs BO = (8 be 4 abe) BO = 2
Slide 134
N has 5 valence electrons 2 N = 10e () = 1e total = 11e Example
10.10: Draw a molecular orbital diagram of N 2 ion and predict its
bond order and magnetic properties Write a MO diagram for N 2 using
N 2 as a base Count the number of valence electrons and assign
these to the MOs following the aufbau principle, Pauli principle
& Hunds rule 2s2s 2s2s 2p2p 2p2p 2p2p 2p2p
Slide 135
Example 10.10: Draw a molecular orbital diagram of N 2 ion and
predict its bond order and magnetic properties Calculate the bond
order by taking the number of bonding electrons and subtracting the
number of antibonding electrons, then dividing by 2 Determine
whether the ion is paramagnetic or diamagnetic 2s2s 2s2s 2p2p 2p2p
2p2p 2p2p BO = (8 be 3 abe) BO = 2.5 Because this is lower than the
bond order in N 2, the bond should be weaker Because there are
unpaired electrons, this ion is paramagnetic
Slide 136
Practice Draw a molecular orbital diagram of C 2 + and predict
its bond order and magnetic properties
Slide 137
2s2s 2s2s 2p2p 2p2p 2p2p 2p2p BO = (5 be 2 abe) BO = 1.5
Because there are unpaired electrons, this ion is paramagnetic
Practice Draw a molecular orbital diagram of C 2 + and predict its
bond order and magnetic properties C has 4 valence electrons 2 C =
8e (+) = 1e total = 7e
Slide 138
Heteronuclear Diatomic Molecules & Ions If the combining
atomic orbitals are identical and equal energy the contribution of
each atomic orbital to the molecular orbital is equal If the
combining atomic orbitals are different types and energies the
atomic orbital closest in energy to the molecular orbital
contributes more to the molecular orbital
Slide 139
Heteronuclear Diatomic Molecules & Ions The more
electronegative an atom is, the lower in energy are its orbitals
Lower energy atomic orbitals contribute more to the bonding MOs
Higher energy atomic orbitals contribute more to the antibonding
MOs Nonbonding MOs remain localized on the atom donating its atomic
orbitals
Slide 140
NO a free radical 2s Bonding MO shows more electron density
near O because it is mostly Os 2s atomic orbital BO = (6 be 1 abe)
BO = 2.5
Slide 141
HF
Slide 142
Polyatomic Molecules When many atoms are combined together, the
atomic orbitals of all the atoms are combined to make a set of
molecular orbitals, which are delocalized over the entire molecule
Gives results that better match real molecule properties than
either Lewis or valence bond theories
Slide 143
Ozone, O 3 MO Theory: Delocalized bonding orbital of O 3