Chapter 10 Chemical Bonding II

Chapter 10Chemical Bonding II

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

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

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

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

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

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

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

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

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

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°

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°

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°

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°

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°

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

Not Quite Perfect Geometry

Because the bonds and atom sizes are not identical in formaldehyde, the observed angles are slightly different from ideal

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

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

Bond Angle Distortion from Lone Pairs

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

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°

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

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

Replacing Atoms with Lone Pairsin the Trigonal Bipyramid System

Seesaw Shape

T–Shape

Linear Shape

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°

Square Pyramidal Shape

Square Planar Shape

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

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

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°

Practice – Predict the molecular geometry and bond angles in SiF5

−

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

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

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

SF6 S

F

FF

F F

F

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

Describing the Geometryof Methanol

Describing the Geometryof Glycine

Practice – Predict the molecular geometries in H3BO3

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

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

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. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule.

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

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


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


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.

Practice – Decide whether the following molecules are polarENO = 3.5N = 3.0Cl = 3.0S = 2.5

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

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

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

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

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

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 – 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

Unhybridized C Orbitals Predict the Wrong Bonding & Geometry

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

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

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

Carbon HybridizationsUnhybridized

2s 2p

sp hybridized

2sp

sp2 hybridized

2p

sp3 hybridized

2p

2sp2

2sp3

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

Orbital Diagram of the sp3 Hybridization of C

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

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

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

Methane Formation with sp3 C

Ammonia Formation with sp3 N

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

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

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

Formaldehyde, CH2O Orbital Diagram

sp2 C

sp2 O

1s H

s

ss

1s H

p C p Op

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

sp2 Hybridized AtomsOrbital Diagrams

Unhybridized atom

2s 2p

sp2 hybridized atom

2sp2

2p

C3 s1 p

2s 2p

2sp2

2p

N2 s1 p

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

Hybrid orbitals overlap to form a s bond.Unhybridized p orbitals overlap to form a p bond.

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

･･

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

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

sp Hybridized AtomsOrbital Diagrams

Unhybridized atom

2s 2p

sp hybridized atom

2sp

2p

C2s2p

2s 2p

2sp

2p

N1s2p

HCN Orbital Diagram

sp C

sp N

1s H

s

s

p C p N2p

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

sp3d Hybridized AtomsOrbital DiagramsUnhybridized atom

3s 3p

sp3d hybridized atom

3sp3d

S3s 3p

3sp3d

P

3d

3d

(non-hybridizing d orbitals not shown)

SOF4 Orbital Diagram

sp3d S

sp2 O

2p F

s

s

d S p Op

2p F

s

2p F

s

2p F

s

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

sp3d2 Hybridized AtomsOrbital Diagrams

Unhybridized atom sp3d2 hybridized atom

S3s 3p 3sp3d2

↑↓

3d

↑↓ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑

I5s 5p↑↓

5d

↑↓ ↑↓ ↑

5sp3d2

↑↓ ↑ ↑ ↑ ↑ ↑

(non-hybridizing d orbitals not shown)

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

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


Determine the hybridization of the interior atoms

C1 = tetrahedral C1 = sp3

C2 = trigonal planar C2 = sp2

Sketch the molecule and orbitals

Label the bonds


Practice – Predict the hybridization of all the atoms in H3BO3

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

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

↑↓

↑↓↑↓

↑↓

↑↓

↑↓

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

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

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

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

Interaction of 1s Orbitals

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

Energy Comparisons of Atomic Orbitals to Molecular Orbitals

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

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

H2

s* Antibonding MOLUMO

s bonding MOHOMO

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

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

s bonding MOHOMO

s* Antibonding MOLUMO

Li2

Interaction of p Orbitals

O2

Dioxygen is paramagneticParamagnetic material has unpaired electronsNeither Lewis theory nor valence bond theory predict this

result

O2 as Described by Lewis and VB Theory

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

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

↑↓

↑ ↑↓ ↓

↑↓

↑

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

Practice – Draw a molecular orbital diagram of C2+

and predict its bond order and magnetic properties

↑↓ 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−

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

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

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

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

Ozone, O3

MO Theory: Delocalized p bonding orbital of O3

Chapter 10 Chemical Bonding II

Documents

Transcript of Chapter 10 Chemical Bonding II