Structure Determines Properties!

Tro, Chemistry: A Molecular Approach 1

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

Chapter 10 Chemical Bonding II


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


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 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


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


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


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 ••

••

••

••

••••

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


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

52

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


Interaction of p Orbitals


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


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


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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Structure Determines Properties!