Chapter 10 Chemical Bonding II
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Transcript of Chapter 10 Chemical Bonding II
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Chapter 10Chemical Bonding II
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TasteThe taste of a food depends on the interaction between the
food molecules and taste cells on your tongueThe 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 cellsWhen this happens, changes in the protein structure cause
a nerve signal to transmit
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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 sugarmaking them “sweeter” than sugar
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Structure Determines Properties!Properties of molecular substances depend on the
structure of the moleculeThe 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
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Molecular GeometryMolecules are 3-dimensional objectsWe often describe the shape of a molecule with terms
that relate to geometric figuresThese 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
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Lewis Theory PredictsElectron 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
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Using Lewis Theory to PredictMolecular Shapes
Lewis theory says that these regions of electron groups should repel each otherbecause 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
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VSEPR TheoryElectron 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 possibleThe resulting geometric arrangement will allow us to predict
the shapes and bond angles in the molecule
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Electron GroupsThe 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
O N O ••••••
•••••• there are three electron groups on N
three lone pairone single bondone double bond
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Electron Group GeometryThere are five basic arrangements of electron groups around a central atombased on a maximum of six bonding electron groups though there may be more than six on very large atoms, it is very
rareEach 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 doesn’t matter which resonance form you use – the electron geometry will be the same
Tro: Chemistry: A Molecular Approach, 2/e
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LinearWhen 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°
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Trigonal PlanarWhen 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°
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TetrahedralWhen 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°
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Trigonal Bipyramidal5 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°
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OctahedralWhen 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 sidesAll positions are equivalentThe bond angle is 90°
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Molecular GeometryThe actual geometry of the molecule may be different from
the electron geometryWhen 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
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Not Quite Perfect Geometry
Because the bonds and atom sizes are not identical in formaldehyde, the observed angles are slightly different from ideal
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The Effect of Lone PairsLone pair groups “occupy more space” on the central atombecause their electron density is exclusively on the
central atom rather than shared like bonding electron groups
Relative sizes of repulsive force interactions isLone 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
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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
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Bond Angle Distortion from Lone Pairs
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Bond Angle Distortion from Lone Pairs
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Bent Molecular Geometry:Derivative of Trigonal Planar Electron GeometryWhen 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
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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 apexBond angle is less than 109.5°
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Bent Tetrahedral Molecular Geometry: A Derivative of Tetrahedral Electron GeometryWhen 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 planar—bent shape, except the angles are
smaller the bond angle is less than 109.5°
Tro: Chemistry: A Molecular Approach, 2/e
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Derivatives of theTrigonal 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° axial–to–axial
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Replacing Atoms with Lone Pairsin the Trigonal Bipyramid System
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Seesaw Shape
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T–Shape
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T–Shape
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Linear Shape
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Derivatives of the Octahedral GeometryWhen 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°
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Square Pyramidal Shape
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Square Planar Shape
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Predicting the Shapes Around Central Atoms
1. Draw the Lewis structure2. Determine the number of electron groups
around the central atom3. 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
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Example 10.2: Predict the geometry and bond angles of PCl3
1. Draw the Lewis structure
a) 26 valence electrons2. Determine the Number
of electron groups around central atom
a) four electron groups around P
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Example 10.2: Predict the geometry and bond angles of PCl3
3. Classify the electron groupsa) three bonding groupsb) 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°
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Practice – Predict the molecular geometry and bond angles in SiF5
−
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Practice – Predict the molecular geometry and bond angles in SiF5
─
Si = 4e─
F5 = 5(7e─) = 35e─
(─) = 1e─
total = 40e─
5 electron groups on Si
5 bonding groups0 lone pairs
Shape = trigonal bipyramid
Bond anglesFeq–Si–Feq = 120°Feq–Si–Fax = 90°
Si least electronegative
Si is central atom
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Practice – Predict the molecular geometry and bond angles in ClO2F
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Practice – Predict the molecular geometry and bond angles in ClO2F
Cl = 7e─
O2 = 2(6e─) = 12e─
F = 7e─
Total = 26e─
4 electron groups on Cl
3 bonding groups1 lone pair
Shape = trigonal pyramidal
Bond anglesO–Cl–O < 109.5°O–Cl–F < 109.5°
Cl least electronegative
Cl is central atom
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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 wedgeFor atoms behind the plane, use a hashed wedge
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SF6 S
F
FF
F F
F
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Multiple Central AtomsMany 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
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Describing the Geometryof Methanol
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Describing the Geometryof Glycine
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Practice – Predict the molecular geometries in H3BO3
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3 electron groups on B
B has3 bonding groups0 pone pairs
4 electron groups on O
O has2 bonding groups2 lone pairs
Practice – Predict the molecular geometries in H3BO3
B = 3e─
O3 = 3(6e─) = 18e─
H3 = 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
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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
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Molecule Polarity
The H─Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule.
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Vector Addition
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Molecule Polarity
The O─C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule.
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Molecule Polarity
The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule.
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Predicting Polarity of Molecules1. Draw the Lewis structure and determine the molecular
geometry2. Determine whether the bonds in the molecule are polar
a) if there are not polar bonds, the molecule is nonpolar3. Determine whether the polar bonds add together to give a
net dipole moment
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Example 10.5: Predict whether NH3 is a polar molecule
1. Draw the Lewis structure and determine the molecular geometry
a) eight valence electronsb) three bonding + one lone pair =
trigonal pyramidal molecular geometry
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Example 10.5: Predict whether NH3 is a polar molecule
2. Determine if the bonds are polar
a) electronegativity differenceb) if the bonds are not polar, we can
stop here and declare the molecule will be nonpolar ENN = 3.0
ENH = 2.13.0 − 2.1 = 0.9therefore the bonds are polar covalent
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Example 10.5: Predict whether NH3 is a polar molecule
3) Determine whether the polar bonds add together to give a net dipole moment
a) vector additionb) generally, asymmetric
shapes result in uncompensated polarities and a net dipole moment
The H─N 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.
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Practice – Decide whether the following molecules are polarENO = 3.5N = 3.0Cl = 3.0S = 2.5
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Practice – Decide whether the following molecules Are polar
polarnonpolar
1. polar bonds, N-O2. asymmetrical shape 1. polar bonds, all S-O
2. symmetrical shape
TrigonalBent Trigonal
Planar2.5
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Molecular Polarity Affects Solubility in WaterPolar molecules are attracted to other
polar moleculesBecause 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
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Problems with Lewis TheoryLewis 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. O2 is paramagnetic, though the Lewis structure predicts it is
diamagnetic
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Valence Bond TheoryLinus 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
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Orbital InteractionAs 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 moleculeThe 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
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Orbital Diagram for the Formation of H2S
+
↑
H
1s ↑↓ H─S bond
↑
H
1sS↑ ↑ ↑↓↑↓
3s 3p
↑↓ H─S bond
Predicts bond angle = 90°Actual bond angle = 92°
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Valence Bond Theory – HybridizationOne 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 = 2s22px
12py12pz
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 placeone hybridization of C is to mix all the 2s and 2p orbitals
to get four orbitals that point at the corners of a tetrahedron
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Unhybridized C Orbitals Predict the Wrong Bonding & Geometry
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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 orbitala) the electrons must be spin paired
3. The shape of the molecule is determined by the geometry of the interacting orbitals
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HybridizationSome 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, sp2, sp3, sp3d, sp3d2
Same type of atom can have different types of hybridization C = sp, sp2, sp3
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Hybrid OrbitalsThe 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
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Carbon HybridizationsUnhybridized
2s 2p
sp hybridized
2sp
sp2 hybridized
2p
sp3 hybridized
2p
2sp2
2sp3
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sp3 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
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Orbital Diagram of the sp3 Hybridization of C
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sp3 Hybridized AtomsOrbital Diagrams
Unhybridized atom
2s 2p
sp3 hybridized atom
2sp3
C
2s 2p
2sp3
N
• Place electrons into hybrid and unhybridized valence orbitals
as if all the orbitals have equal energy• Lone pairs generally occupy hybrid orbitals
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Practice – Draw the orbital diagram for the sp3 hybridization of each atom
Unhybridized atom
3s 3p Cl
2s 2p O
sp3 hybridized atom
3sp3
2sp3
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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
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Methane Formation with sp3 C
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Ammonia Formation with sp3 N
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Types of BondsA sigma (s) bond results when the interacting atomic orbitals
point along the axis connecting the two bonding nuclei either standard atomic orbitals or hybrids
s–to–s, p–to–p, hybrid–to–hybrid, s–to–hybrid, etc.
A pi (p) 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 s bonds are stronger than p bonds
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Orbital Diagrams of Bonding“Overlap” between a hybrid orbital on one atom with a hybrid or nonhybridized orbital on another atom results in a s bond
“Overlap” between unhybridized p orbitals on bonded atoms results in a p bond
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CH3NH2 Orbital Diagram
sp3 C sp3 N
1s H
s
sss s s
1s H 1s H 1s H 1s H
s s
s ss
s
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Formaldehyde, CH2O Orbital Diagram
sp2 C
sp2 O
1s H
s
ss
1s H
p C p Op
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sp2
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 s bonds and lone pairs, uses nonhybridized p orbital for p bond
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sp2 Hybridized AtomsOrbital Diagrams
Unhybridized atom
2s 2p
sp2 hybridized atom
2sp2
2p
C3 s1 p
2s 2p
2sp2
2p
N2 s1 p
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Practice – Draw the orbital diagram for the sp2 hybridization of each atom. How many s and p bonds would you expect each to form?
Unhybridized atom
2s 2p
B3 s0 p
2s 2p
O1 s1 p
sp2 hybridized atom
2sp2
2p
2sp2
2p
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Hybrid orbitals overlap to form a s bond.Unhybridized p orbitals overlap to form a p bond.
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CH2NH Orbital Diagram
sp2 C
sp2 N
1s H
s
ss s
1s H 1s H
p C p Np
C N
H
H H
・・
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Bond RotationOrbitals form the s bond point along the internuclear axis
rotation around that bond does not require breaking the interaction between the orbitals
Orbitals that form the p bond interact above and below the internuclear axis rotation around the axis requires the breaking of the
interaction between the orbitals
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spAtom with two electron groups
linear shape180° bond angle
Atom uses hybrid orbitals for s bonds or lone pairs, uses nonhybridized p orbitals for p bonds
p
ps
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sp Hybridized AtomsOrbital Diagrams
Unhybridized atom
2s 2p
sp hybridized atom
2sp
2p
C2s2p
2s 2p
2sp
2p
N1s2p
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HCN Orbital Diagram
sp C
sp N
1s H
s
s
p C p N2p
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sp3d Atom with five electron groups around it
trigonal bipyramid electron geometry Seesaw, T–Shape, Linear 120° & 90° bond angles
Use empty d orbitals from valence shell d orbitals can be used to make p bonds
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sp3d Hybridized AtomsOrbital DiagramsUnhybridized atom
3s 3p
sp3d hybridized atom
3sp3d
S3s 3p
3sp3d
P
3d
3d
(non-hybridizing d orbitals not shown)
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SOF4 Orbital Diagram
sp3d S
sp2 O
2p F
s
s
d S p Op
2p F
s
2p F
s
2p F
s
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sp3d2 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 p bonds
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sp3d2 Hybridized AtomsOrbital Diagrams
Unhybridized atom sp3d2 hybridized atom
S3s 3p 3sp3d2
↑↓
3d
↑↓ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑
I5s 5p↑↓
5d
↑↓ ↑↓ ↑
5sp3d2
↑↓ ↑ ↑ ↑ ↑ ↑
(non-hybridizing d orbitals not shown)
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Predicting Hybridization and Bonding Scheme
1. Start by drawing the Lewis structure2. Use VSEPR Theory to predict the electron
group geometry around each central atom3. 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 s or p
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Example 10.7: Predict the hybridization and bonding scheme for CH3CHO
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
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Example 10.7: Predict the hybridization and bonding scheme for CH3CHO
Determine the hybridization of the interior atoms
C1 = tetrahedral C1 = sp3
C2 = trigonal planar C2 = sp2
Sketch the molecule and orbitals
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Label the bonds
Example 10.7: Predict the hybridization and bonding scheme for CH3CHO
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Practice – Predict the hybridization of all the atoms in H3BO3
H = can’t hybridizeB = 3 electron groups = sp2
O = 4 electron groups = sp3
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Practice – Predict the hybridization and bonding scheme of all the atoms in NClO
N = 3 electron groups = sp2
O = 3 electron groups = sp2
Cl = 4 electron groups = sp3
••O N Cl ••••
••••••
s:Nsp2─Clp
s:Osp2─Nsp2
p:Op─Np
↑↓
↑↓↑↓
↑↓
↑↓
↑↓
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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 doesn’t predict perfectly magnetic behavior of O2
In addition, VB theory presumes the electrons are localized in orbitals on the atoms in the molecule – it doesn’t account for delocalization
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Molecular Orbital TheoryIn MO theory, we apply Schrödinger’s 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 minimizedIn this treatment, the electrons belong to the whole molecule – so the orbitals belong to the whole moleculedelocalization
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LCAOThe 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 methodweighted sum
Because the orbitals are wave functions, the waves can combine either constructively or destructively
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Molecular OrbitalsWhen the wave functions combine constructively, the resulting molecular orbital has less energy than the original atomic orbitals – it is called a Bonding Molecular Orbitals, pmost 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 Orbitals*, p*most of the electron density outside the nuclei nodes between nuclei
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Interaction of 1s Orbitals
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Molecular Orbital TheoryElectrons in bonding MOs are stabilizing
lower energy than the atomic orbitalsElectrons in antibonding MOs are destabilizinghigher in energy than atomic orbitalselectron density located outside the internuclear axis
electrons in antibonding orbitals cancel stability gained by electrons in bonding orbitals
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Energy Comparisons of Atomic Orbitals to Molecular Orbitals
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MO and PropertiesBond Order = difference between number of electrons in bonding and antibonding orbitalsonly need to consider valence electronsmay be a fractionhigher bond order = stronger and shorter bonds if bond order = 0, then bond is unstable compared to
individual atoms and no bond will formA substance will be paramagnetic if its MO diagram has unpaired electrons if all electrons paired it is diamagnetic
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1s 1s
s*
s
Hydrogen AtomicOrbital
Hydrogen AtomicOrbital
Dihydrogen, H2 MolecularOrbitals
Because more electrons are in bonding orbitals than are in antibonding orbitals,
net bonding interaction
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H2
s* Antibonding MOLUMO
s bonding MOHOMO
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1s 1s
s*
s
Helium AtomicOrbital
Helium AtomicOrbital
Dihelium, He2 MolecularOrbitals
Because there are as many electrons in antibonding orbitals as in bonding orbitals,
there is no net bonding interaction
BO = ½(2-2) = 0
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1s 1s
s*
s
Lithium AtomicOrbitals
Lithium AtomicOrbitals
Dilithium, Li2 MolecularOrbitals
Because more electrons are in bonding orbitals than are in antibonding orbitals, there is a
net bonding interaction
2s 2s
s*
s Any fill energy level will generate filled bonding and antibonding MO’s;therefore only need to consider valence shell
BO = ½(4-2) = 1
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s bonding MOHOMO
s* Antibonding MOLUMO
Li2
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Interaction of p Orbitals
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Interaction of p Orbitals
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O2
Dioxygen is paramagneticParamagnetic material has unpaired electronsNeither Lewis theory nor valence bond theory predict this
result
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O2 as Described by Lewis and VB Theory
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OxygenAtomicOrbitals
OxygenAtomicOrbitals
2s 2s
2p 2p
s*
s
s*
s
p*
pBecause 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, O2 is predicted to be
paramagnetic
O2 MO’s
BO = ½(8 be – 4 abe)BO = 2
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N has 5 valence electrons
2 N = 10e−
(−) = 1e−
total = 11e−
Example 10.10: Draw a molecular orbital diagram of N2− ion
and predict its bond order and magnetic properties
Write a MO diagram for N2
− using N2 as a base
Count the number of valence electrons and assign these to the MOs following the aufbau principle, Pauli principle & Hund’s rule
↑↓ s2s
s*2s
p2p
s2p
p*2p
s*2p
↑↓
↑ ↑↓ ↓
↑↓
↑
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Example 10.10: Draw a molecular orbital diagram of N2− ion
and predict its bond order and magnetic propertiesCalculate the bond order by taking the number of bonding electrons and subtracting the number of antibonding electrons, then dividing by 2Determine whether the ion is paramagnetic or diamagnetic
↑↓ s2s
s*2s
p2p
s2p
p*2p
s*2p
↑↓
↑ ↑↓ ↓
↑↓
↑
BO = ½(8 be – 3 abe)BO = 2.5
Because this is lower than the bond order
in N2, the bond should be weaker
Because there are unpaired electrons, this ion is paramagnetic
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Practice – Draw a molecular orbital diagram of C2+
and predict its bond order and magnetic properties
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↑↓ s2s
s*2s
p2p
s2p
p*2p
s*2p
↑↓
↑ ↑↓BO = ½(5 be – 2 abe)
BO = 1.5
Because there are unpaired electrons, this ion is paramagnetic
Practice – Draw a molecular orbital diagram of C2+
and predict its bond order and magnetic propertiesC has 4 valence electrons
2 C = 8e−
(+) = −1e−
total = 7e−
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Heteronuclear Diatomic Molecules & IonsIf 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
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Heteronuclear Diatomic Molecules & IonsThe 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
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NO
a free radical
s2s Bonding MOshows more
electron density near O because it is mostly O’s
2s atomic orbital
BO = ½(6 be – 1 abe)BO = 2.5
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HF
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Polyatomic MoleculesWhen 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
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Ozone, O3
MO Theory: Delocalized p bonding orbital of O3