chem chapter10 lec

Chapter 10Chemical

Bonding II

2008, Prentice Hall

Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro

Roy KennedyMassachusetts Bay Community College

Wellesley Hills, MA

Tro, Chemistry: A Molecular Approach 2

Structure Determines Properties!• properties of molecular substances depend on

the structure of the molecule• the structure includes many factors, including:

the skeletal arrangement of the atomsthe kind of bonding between the atoms

ionic, polar covalent, or covalentthe shape of the molecule

• bonding theory should allow you to predict the shapes of molecules


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


Using Lewis Theory to PredictMolecular Shapes

• Lewis theory predicts there are regions of electrons in an atom based on placing shared pairs of valence electrons between bonding nuclei and unshared valence electrons located on single nuclei

• this idea can then be extended to predict the shapes of molecules by realizing these regions are all negatively charged and should repel


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 theorysince 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


Electron Groups• the Lewis structure predicts the arrangement of valence

electrons 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 3 electron groups on N1 lone pair1 single bond1 double bond


Molecular Geometries• there are 5 basic arrangements of electron groups

around a central atombased on a maximum of 6 bonding electron groups

though there may be more than 6 on very large atoms, it is very rare

• each of these 5 basic arrangements results in 5 different basic molecular shapes 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 molecular geometry will be the same


Linear Geometry• when there are 2 electron groups around the central

atom, they will occupy positions opposite each other around the central atom

• this results in the molecule taking a linear geometry• the bond angle is 180°

ClBeCl

O C O


Linear Geometry


Trigonal Geometry• when there are 3 electron groups around the central

atom, they will occupy positions in the shape of a triangle around the central atom

• this results in the molecule taking a trigonal planar geometry

• the bond angle is 120°

F

F B F


Trigonal Geometry


Not Quite Perfect Geometry

Because the bonds are not identical, the observed angles are slightly different from ideal.


Tetrahedral Geometry• when there are 4 electron groups around the central

atom, they will occupy positions in the shape of a tetrahedron around the central atom

• this results in the molecule taking a tetrahedral geometry

• the bond angle is 109.5°

F

F C F

F


Tetrahedral Geometry


Methane


Trigonal Bipyramidal Geometry• when there are 5 electron groups around the central atom, they

will occupy positions in the shape of a two tetrahedra that are base-to-base with the central atom in the center of the shared bases

• this results in the molecule taking a trigonal bipyramidal geometry

• 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°


Trigonal Bipyramid


Trigonal Bipyramidal Geometry

P

ClCl

ClCl

Cl••

•• •

•••

•••

•

••••

••

••

••••

••

••

••


Octahedral Geometry• when there are 6 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 molecule taking an octahedral geometry it is called octahedral because the geometric figure has 8

sides

• all positions are equivalent• the bond angle is 90°


Octahedral Geometry


Octahedral Geometry

S

F F

FF

F

F

••

•• •

•••

••

••

••

••

••••

••

••

••••

••

••

••••


The Effect of Lone Pairs• lone 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 is:Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair • this effects the bond angles, making them smaller

than expected


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.


Derivative Shapes

• the molecule’s shape will be one of basic molecular geometries if all the electron groups are bonds and all the bonds are equivalent

• molecules with lone pairs or different kinds of surrounding atoms will have distorted bond angles and different bond lengths, but the shape will be a derivative of one of the basic shapes


Derivative of Trigonal Geometry• when there are 3 electron groups around the central

atom, and 1 of them is a lone pair, the resulting shape of the molecule is called a trigonal planar - bent shape

• the bond angle is < 120°

O S O

O S O

O S O


Derivatives of Tetrahedral Geometry

• when there are 4 electron groups around the central atom, and 1 is a lone pair, the result is called a pyramidal shapebecause it is a triangular-base pyramid with the central

atom at the apex

• when there are 4 electron groups around the central atom, and 2 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

• for both shapes, the bond angle is < 109.5°


Pyramidal Shape


Bond Angle Distortion from Lone Pairs


Tetrahedral-Bent Shape


Pyramidal Shape

F

FAsF


Bond Angle Distortion from Lone Pairs


Tetrahedral-Bent Shape

1

OClO


Derivatives of theTrigonal Bipyramidal Geometry

• when there are 5 electron groups around the central atom, and some are lone pairs, they will occupy the equatorial positions because there is more room

• when there are 5 electron groups around the central atom, and 1 is a lone pair, the result is called see-saw shape aka distorted tetrahedron

• when there are 5 electron groups around the central atom, and 2 are lone pairs, the result is called T-shaped

• when there are 5 electron groups around the central atom, and 3 are lone pairs, the result is called a linear shape

• the bond angles between equatorial positions is < 120°

• the bond angles between axial and equatorial positions is < 90° linear = 180° axial-to-axial


Replacing Atoms with Lone Pairsin the Trigonal Bipyramid System


See-Saw Shape

F S F

F

F

•••

•

••

••

••

••

••

••

••

••

••

••

••


T-Shape


T-Shaped

F Cl F

F

•••

•

••

••

•• ••

••

••

••

••

••


Linear Shape


Derivatives of theOctahedral Geometry

• when there are 6 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

• when there are 6 electron groups around the central atom, and 1 is a lone pair, the result is called a square pyramid shape the bond angles between axial and equatorial positions is < 90°

• when there are 6 electron groups around the central atom, and 2 are lone pairs, the result is called a square planar shape the bond angles between equatorial positions is 90°


Square Pyramidal Shape

Br

FF

FF

F••

•• •

•••

•••

•

••••

••

••

••••

••

••

••

••


Square Planar Shape

F Xe F

F

F

•••

•

••

••

••

••

••

••

••

••

••

••

••

••


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 1 group

4) Use Table 10.1 to Determine the Shape and Bond Angles


Practice – Predict the Molecular Geometry and Bond Angles in SiF5

-1


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

Si

FF

FF

F••

•• •

•••

•••

•

••••

••

••

••••

••

••

••

-1

Bond AnglesFeq-Si-Feq = 120°Feq-Si-Fax = 90°

Si Least Electronegative

Si Is Central Atom


Practice – Predict the Molecular Geometry and Bond Angles in ClO2F


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

O Cl

O

F

••

••

•• •

•••

••

••

••

••

••


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


SF6

S

F

F

F

F F

F

S

F F

FF

F

F

••

•• •

•••

••

••

••

••

••••

••

••

••••

••

••

••••


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

H|

HOCCH|||

OH

shape around left C is tetrahedral

shape around center C is trigonal planar

shape around right O is tetrahedral-bent


Describing the Geometryof Methanol


Describing the Geometryof Glycine


Practice – Predict the Molecular Geometries in H3BO3

59

Practice – Predict the Molecular Geometries in H3BO3

B = 3e─

O3 = 3(6e─) = 18e─

H3 = 3(1e─) = 3e─

Total = 24e─

3 Electron Groups on B

B has3 Bonding Groups0 Lone Pairs

Shape on B = Trigonal Planar

B Least Electronegative

B Is Central Atom

oxyacid, so H attached to O

O B

O

OH H

H••

••

••

••

••

•• 4 Electron Groups on O

O has2 Bonding Groups2 Lone Pairs

Shape on O = Tetrahedral Bent


Polarity of Molecules

• in order for a molecule to be polar it must1) 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


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.


Vector Addition


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.


Molecule Polarity

The H-O bond is polar. The both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule.


Molecule Polarity

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.


Molecular Polarity Affects

Solubility in Water• polar molecules are attracted to

other polar molecules• since water is a polar molecule,

other polar molecules dissolve well in waterand ionic compounds as well

• some molecules have both polar and nonpolar parts


A Soap MoleculeSodium Stearate


Practice - Decide Whether the Following Are Polar

O N Cl ••

••

••

••

••••

O S

O

O

••

••

•• •

•••••

••

••ENO = 3.5N = 3.0Cl = 3.0S = 2.5


Practice - Decide Whether the Following Are Polar

polarnonpolar

1) polar bonds, N-O2) asymmetrical shape 1) polar bonds, all S-O

2) symmetrical shape

O N Cl ••

••

••

••

••••

O S

O

O

••

••

•• •

•••••

••

••

TrigonalBent Trigonal

PlanarCl

N

O

3.0

3.0

3.5

O

O

OS

3.5

3.5 3.52.5


Problems with Lewis Theory• 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 moleculese.g., O2 is paramagnetic, though the Lewis structure

predicts it is diamagnetic


Valence Bond Theory

• Linus Pauling and others applied the principles of quantum mechanics to molecules

• they reasoned that bonds between atoms would arise 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


Orbital Interaction• as two atoms approached, the partially filled or

empty valence atomic orbitals on the atoms would interact to form molecular orbitals

• the molecular orbitals would be more stable than the separate atomic orbitals because they would contain paired electrons shared by both atomsthe interaction energy between atomic orbitals is

negative when the interacting atomic orbitals contain a total of 2 electrons


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°


Valence Bond Theory - Hybridization• one of the issues that arose was that the number of

partially filled or empty atomic orbital did not predict the number of bonds or orientation of bonds C = 2s22px

12py12pz

0 would predict 2 or 3 bonds that are 90° apart, rather than 4 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 4 orbitals that point at the corners of a tetrahedron


Unhybridized C Orbitals Predict the Wrong Bonding & Geometry


Valence Bond TheoryMain Concepts

1. the valence electrons in an atom reside in the quantum mechanical atomic orbitals or hybrid orbitals

2. a chemical bond results when these atomic orbitals overlap and there is a total of 2 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 overlapping orbitals


Hybridization• some atoms hybridize their orbitals to maximize

bondinghybridizing is mixing different types of orbitals to

make a new set of degenerate orbitalssp, sp2, sp3, sp3d, sp3d2

more bonds = more full orbitals = more stability

• better explain observed shapes of molecules

• same type of atom can have different hybridization depending on the compoundC = sp, sp2, sp3


Hybrid Orbitals• H cannot hybridize!!

• the number of standard atomic orbitals combined = the number of hybrid orbitals formed

• 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 in other words, you have to know the structure of the

molecule beforehand in order to predict the hybridization


Orbital Diagrams with Hybridization

• place electrons into hybrid and unhybridized valence orbitals as if all the orbitals have equal energy

• when bonding, bonds form between hybrid orbitals and bonds form between unhybridized orbitals that are parallel


Carbon Hybridizations

Unhybridized

2s 2p

sp hybridized

2sp

sp2 hybridized

2p

sp3 hybridized

2p

2sp2

2sp3


sp3 Hybridization• atom with 4 areas of electrons

tetrahedral geometry109.5° angles between hybrid orbitals

• atom uses hybrid orbitals for all bonds and lone pairs

H C N H

H

H H

s

•• sp3

s

sp3


sp3 Hybridization of C


sp3 Hybridized AtomsOrbital Diagrams

Unhybridized atom

2s 2p

sp3 hybridized atom

2sp3

C

2s 2p

2sp3

N


Methane Formation with sp3 C


Ammonia Formation with sp3 N


CH3NH2 Orbital Diagram

sp3 C sp3 N

1s H

1s H 1s H 1s H 1s H


Practice - Draw the Orbital Diagram for the sp3 Hybridization of Each Atom

Unhybridized atom

3s 3p

Cl

2s 2p

O


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


Types of Bonds• a sigma () bond results when the bonding atomic

orbitals point along the axis connecting the two bonding nucleieither standard atomic orbitals or hybrids

s-to-s, p-to-p, hybrid-to-hybrid, s-to-hybrid, etc.

• a pi () bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nucleibetween 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


Bond Rotation• because orbitals that form the bond point

along the internuclear axis, rotation around that bond does not require breaking the interaction between the orbitals

• but the orbitals that form the bond interact above and below the internuclear axis, so rotation around the axis requires the breaking of the interaction between the orbitals


sp2

• atom with 3 areas of electrons trigonal planar system

C = trigonal planar N = trigonal bent O = “linear”

120° bond anglesflat

• atom uses hybrid orbitals for bonds and lone pairs, uses nonhybridized p orbital for bond

H C O H

O ••

sp2s ••

••

••sp2

sp3 s


+

sp2 sp2

Hybrid orbitals overlap to form bondUnhybridized p orbitals overlap to form bond


sp2 Hybridized AtomsOrbital Diagrams

Unhybridized atom

2s 2p

sp2 hybridized atom

2sp2

2p

C3 1

2s 2p

2sp2

2p

N2 1


CH2NH Orbital Diagram

sp2 C

sp2 N

1s H

1s H 1s H

p C p N


Practice - Draw the Orbital Diagram for the sp2 Hybridization of Each Atom. How many and bonds would you expect each to form?

Unhybridized atom

2s 2p

B

2s 2p

O


Practice - Draw the Orbital Diagram for the sp2 Hybridization of Each Atom. How many and bonds would you expect each to form?

Unhybridized atom

2s 2p

B

3 0

2s 2p

O1 1

sp2 hybridized atom

2sp2

2p

2sp2

2p


sp• atom with 2 areas of electrons

linear shape180° bond angle

• atom uses hybrid orbitals for bonds or lone pairs, uses nonhybridized p orbitals for bonds

H C N

sp sps


sp Hybridized AtomsOrbital Diagrams

Unhybridized atom

2s 2p

sp hybridized atom

2sp

2p

C22

2s 2p

2sp

2p

N12


HCN Orbital Diagram

sp C

sp N

1s H

p C p N


sp3d• atom with 5 areas of electrons

around ittrigonal bipyramid shapeSee-Saw, T-Shape, Linear120° & 90° bond angles

• use empty d orbitals from valence shell

• d orbitals can be used to make bonds

••

I

F

FO

OO

••

••••

••••

••

••

••

••

••

••

••

-1


sp3d Hybridized AtomsOrbital Diagrams

Unhybridized atom

3s 3p

sp3d hybridized atom

3sp3d

S

3s 3p

3sp3d

P

3d

3d

(non-hybridizing d orbitals not shown)


sp3d2

• atom with 6 areas of electrons around itoctahedral shapeSquare Pyramid, Square Planar90° bond angles

• use empty d orbitals from valence shell

• d orbitals can be used to make bonds

•• ••

••••

••Br

F

F

F

F

F••••

••

•• ••••

•• ••

••

••

••


sp3d2 Hybridized AtomsOrbital Diagrams

Unhybridized atom sp3d2 hybridized atom

S

3s 3p 3sp3d2

↑↓

3d

↑↓ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑

I

5s 5p

↑↓

5d

↑↓ ↑↓ ↑

5sp3d2

↑↓ ↑ ↑ ↑ ↑ ↑

(non-hybridizing d orbitals not shown)


Example - Predict the Hybridization of All the Atoms in H3BO3

O B

O

OH H

H••

••

••

••

••

••

H = can’t hybridizeB = 3 electron groups = sp2

O = 4 electron groups = sp3


Practice - Predict the Hybridization and Bonding Scheme of All the Atoms in NClO

O N Cl ••

••

••

••

••••


Practice - Predict the Hybridization and Bonding Scheme of All the Atoms in NClO

O N Cl ••

••

••

••

••••

N = 3 electron groups = sp2

O = 3 electron groups = sp2

Cl = 4 electron groups = sp3


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 geometry4) Sketch the atomic and hybrid orbitals on the atoms in

the molecule, showing overlap of the appropriate orbitals

5) Label the bonds as or


Ex 10.8 – 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

C 1H

H

H

C 2 H

O



Determine the hybridization of the interior atoms

C1 = tetrahedral

C1 = sp3

C2 = trigonal planar

C2 = sp2

Sketch the molecule and orbitals

C 1H

H

H

C 2 H

O



Label the bonds

C 1H

H

H

C 2 H

O


Problems with Valence Bond Theory

• VB theory predicts many properties better than Lewis Theorybonding schemes, bond strengths, bond lengths,

bond rigidity

• however, there are still many properties of molecules it doesn’t predict perfectlymagnetic behavior of O2


Molecular Orbital Theory• in 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 minimized

• in this treatment, the electrons belong to the whole molecule – so the orbitals belong to the whole moleculeunlike VB Theory where the atomic orbitals still exist in the

molecule


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 methodweighted sum

• because the orbitals are wave functions, the waves can combine either constructively or destructively


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 a Antibonding Molecular Orbital*, *most of the electron density outside the nuclei nodes between nuclei


Interaction of 1s Orbitals


Molecular Orbital Theory• Electrons in bonding MOs are stabilizing

Lower energy than the atomic orbitals

• Electrons in anti-bonding MOs are destabilizingHigher in energy than atomic orbitalsElectron density located outside the

internuclear axisElectrons in anti-bonding orbitals cancel

stability gained by electrons in bonding orbitals


MO and Properties• Bond 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 - no bond will form.

• A substance will be paramagnetic if its MO diagram has unpaired electrons if all electrons paired it is diamagnetic

2

Elec. Antibond# - Elec. Bond # Order Bond


1s 1s

Hydrogen AtomicOrbital

Hydrogen AtomicOrbital

Dihydrogen, H2 MolecularOrbitals

Since more electrons are in bonding orbitals than are in antibonding orbitals,

net bonding interaction


H2

* Antibonding MOLUMO

bonding MOHOMO


1s 1s

Helium AtomicOrbital

Helium AtomicOrbital

Dihelium, He2 MolecularOrbitals

Since there are as many electrons in antibonding orbitals as in bonding orbitals,

there is no net bonding interaction

BO = ½(2-2) = 0

134

1s 1s

Lithium AtomicOrbitals

Lithium AtomicOrbitals

Dilithium, Li2 MolecularOrbitals

Since more electrons are in bonding orbitals than are in

antibonding orbitals, net bonding interaction

2s 2s

Any fill energy level will generate filled bonding and

antibonding MO’s;therefore only need to consider valence shell

BO = ½(4-2) = 1


bonding MOHOMO

Antibonding MOLUMO

Li2


Interaction of p Orbitals


O2

• dioxygen is paramagnetic• paramagnetic material have unpaired electrons• neither Lewis Theory nor Valence Bond Theory

predict this result


O2 as described by Lewis and VB theory


OxygenAtomicOrbitals

OxygenAtomicOrbitals

2s 2s

2p 2p

Since more electrons are in bonding orbitals than are in

antibonding orbitals, net bonding interaction

Since there are unpaired electrons in the

antibonding orbitals, O2 is paramagnetic

O2 MO’s

BO = ½( 8 be – 4 abe)BO = 2


O2

Antibonding MOHOMO

Bonding MOHOMO-1

Antibonding MOLUMO+1

Antibonding MOLUMO


Draw a molecular orbital diagram of N2 and predict its bond order and magnetic properties


NitrogenAtomicOrbitals

NitrogenAtomicOrbitals

2s 2s

2p 2p

Since there are no unpaired electrons, N2 is diamagnetic

N2 MO’s

BO = ½( 8 be – 2 abe)BO = 3


Heteronuclear Diatomic Molecules• the more electronegative atom has lower energy orbitals• when the combining atomic orbitals are identical and

equal energy, the weight of each atomic orbital in the molecular orbital are equal

• when the combining atomic orbitals are different kinds and energies, the atomic orbital closest in energy to the molecular orbital contributes more to the molecular orbital lower energy atomic orbitals contribute more to the bonding MOhigher energy atomic orbitals contribute more to the

antibonding MO

• nonbonding MOs remain localized on the atom donating its atomic orbitals


NO

Free-Radical

2s Bonding MOmainly O’s 2s atomic orbital


HF


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


Ozone, O3

Delocalized bonding orbital of O3

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