molecular structure and bonding

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Molecular Structure and Bonding Dr.Christoph Phayao University June 2014

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How can we describe the bonding between atoms forming molecules ? We start from valence bond theory and come to molecular orbital theory.

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Page 1: Molecular structure and bonding

Molecular Structure and Bonding

Dr.Christoph Phayao University June 2014

Page 2: Molecular structure and bonding
Page 3: Molecular structure and bonding

Part 1

What is a chemical bond ?

Page 4: Molecular structure and bonding

Ionic Bond

Normally between a metal and a non-metal: They exchange electrons and become ions (charged atoms) which attract each other by electrostatic force.

A pair of ions does not stay alone but form crystals

Page 5: Molecular structure and bonding

Covalent Bond

Two non-metals share (valence) electrons:

(Remark: Transition metals can form covalent bonds also !)

Page 6: Molecular structure and bonding

Polar Covalent Bond

Two non-metals share electrons unevenly because of electronegativity difference. Electrons are closer to one atom than the other.

This results on partially negative and positive charges on the atoms

Page 7: Molecular structure and bonding

Metallic Bond

Metal atoms share all their valence electrons, which freely move between all atoms which form a network.

Therefore all metals can conduct electricity and look shiny

Page 8: Molecular structure and bonding

Bond Polarity

Page 9: Molecular structure and bonding

Polar Bonds

Uneven sharing of electrons due to differences in Electronegativity

The “pull” an atom has for electrons

Page 10: Molecular structure and bonding
Page 11: Molecular structure and bonding

Electronegativity Trends

Page 12: Molecular structure and bonding

Common Electronegativites

Highest value, set to 4

Page 13: Molecular structure and bonding
Page 14: Molecular structure and bonding

Polar Molecules

Electrons are not equally shared in a bond, which can lead to a dipolmoment of the whole molecule

Page 15: Molecular structure and bonding

Polar Bonds and Geometry

Page 16: Molecular structure and bonding

Which bond type ?

(exception: Transition metals !)

Page 17: Molecular structure and bonding

Electron counting

Page 18: Molecular structure and bonding

Formal Charge Split all bonds in the middle => “real” charge on atoms

(2) Octet Rule Count all bonding electrons for one atom => 8 is most stable

(3) Oxidation Number Give all bonding electrons to the more electronegative atom

Page 19: Molecular structure and bonding

Special Cases

“Extended octet” Especially P and S can use d-orbitals to make more than 3 resp. 2 bonds !

6 VE: Especially common for B and Al !

Page 20: Molecular structure and bonding

Part 2: Valence Bond Theory (VB)

“Valence Electrons are located in bonds and lone pairs”

Page 21: Molecular structure and bonding
Page 22: Molecular structure and bonding

Sigma bonds

Page 23: Molecular structure and bonding

Pi Bonds

Page 24: Molecular structure and bonding

“Resonance”

Write the resonance formula for OZONE ! Does the molecule have a charge ?

Page 25: Molecular structure and bonding
Page 26: Molecular structure and bonding

Important exception: Carbon Monoxide !

Page 27: Molecular structure and bonding

Ho

mew

ork (2

)

Draw Lewis structure(s) and find formal charges (all atoms) and hybridization (central atom) in: 1. NO3 (-) 2. PO4 (3-) 3. CH3 Cl 4. CH2Cl2 5. SO2 6. SO3 7. CO3 (2-) 8. H2O2 9. N2O 10. Cl O2

Page 28: Molecular structure and bonding

***** Break *****

Page 29: Molecular structure and bonding

Prentice Hall © 2003 Chapter 9

• Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding.

• Hybridization is determined by the electron domain geometry.

sp Hybrid Orbitals

• Consider the BeF2 molecule (experimentally known to exist):

Hybrid Orbitals

Page 30: Molecular structure and bonding

Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.

atomic

orbitals

hybrid

orbitals

orbital box diagrams

Page 31: Molecular structure and bonding

Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued).

orbital box diagrams with orbital contours

Page 32: Molecular structure and bonding
Page 33: Molecular structure and bonding

Figure 11.3 The sp2 hybrid orbitals in BF3.

Page 34: Molecular structure and bonding

sp2 and sp3

Hybrid Orbitals

Page 35: Molecular structure and bonding

Figure 11.4 The sp3 hybrid orbitals in CH4.

Page 36: Molecular structure and bonding

Figure 11.5 The sp3 hybrid orbitals in NH3.

Page 37: Molecular structure and bonding

Figure 11.5 continued The sp3 hybrid orbitals in H2O.

Page 38: Molecular structure and bonding

Including d-orbitals

3d orbitals can be filled as well => Al acts as Lewis acid => P and S have “hypervalence”

Page 39: Molecular structure and bonding

Figure 11.6 The sp3d hybrid orbitals in PCl5.

Page 40: Molecular structure and bonding

Figure 11.7 The sp3d2 hybrid orbitals in SF6.

Page 41: Molecular structure and bonding

Acid or Base ?

Compare AlCl3 and PCl3 ? Which acts as acid and which as base – and why ? Why is FeCl3 a strong Lewis acid ?

Page 42: Molecular structure and bonding
Page 43: Molecular structure and bonding

SOLUTION:

PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO.

PLAN: Use the Lewis structures to ascertain the arrangement of groups and

shape at each central atom. Postulate the hybrid orbitals taking note of

the multiple bonds and their orbital overlaps.

H3C

C

CH3

O

sp3 hybridized

sp3 hybridized

CC

C

O

H

H

HHH

H

sp2 hybridized

bonds bond

CC

C

O

sp3

sp3

sp3

sp3

sp3

sp3

sp3

sp3

sp2 sp2

sp2

sp2

sp2sp2

H

HH

HH

H

Page 44: Molecular structure and bonding

Tasks

• Draw the Lewis Structures and the Hybrid Orbitals for Ethane, Ethene and Ethyne (mark the hybrid orbitals)

• Which hybridization has the central atom in: H2O, O2, NH3, NH4+, N in pyridine, O in THF, S in SOCl2, C in HCHO compared to CO

Page 45: Molecular structure and bonding

Chemical Reactivity

From the hybrid orbitals we can estimate if a molecule acts as Lewis acid or base (if there is an electrophilic or nucleophilic center) Consider the “empty” pz orbital of C in HCHO vs. the “filled” sp orbital of C in CO -> in the first case, it acts as Lewis acid, in the second as base !

Page 46: Molecular structure and bonding

***** Break *****

Page 47: Molecular structure and bonding

VSEPR

Page 48: Molecular structure and bonding

VSEPR Theory Clip: http://www.youtube.com/watch?v=nxebQZUVvTg

Page 49: Molecular structure and bonding

MO Theory

Page 50: Molecular structure and bonding

The Central Themes of MO

Theory

A molecule is viewed on a quantum mechanical level as a collection of nuclei

surrounded by delocalized molecular orbitals.

Atomic wave functions are summed to obtain molecular wave functions.

If wave functions reinforce each other, a bonding MO is formed (region of

high electron density exists between the nuclei).

If wave functions cancel each other, an antibonding MO is formed (a node of

zero electron density occurs between the nuclei).

Page 51: Molecular structure and bonding

Amplitudes of wave

functions added

Figure 11.14

An analogy between light waves and atomic wave functions.

Amplitudes of

wave functions

subtracted.

Page 52: Molecular structure and bonding

Prentice Hall © 2003 Chapter 9

Molecular Orbitals

• Molecular orbitals:

• each contain a maximum of two electrons

• have definite energies

• can be visualized with contour diagrams

• are distributed over the whole molecule (not only in between 2 atoms)

• When two AOs overlap, two MOs form.

Page 53: Molecular structure and bonding

Prentice Hall © 2003 Chapter 9

Molecular Orbitals

The Hydrogen Molecule

Page 54: Molecular structure and bonding

Prentice Hall © 2003 Chapter 9

Figure 11.15 The MO diagram for H2.

En

erg

y

MO

of H2

*1s

1s

AO

of H

1s

AO

of H

1s

H2 bond order

= 1/2(2-0) = 1

Filling molecular orbitals with electrons follows the

same concept as filling atomic orbitals.

Page 55: Molecular structure and bonding
Page 56: Molecular structure and bonding

Prentice Hall © 2003 Chapter 9

Electron Configurations and Molecular

Properties

• Two types of magnetic behavior:

• paramagnetism (unpaired electrons in molecule): strong attraction between magnetic field and molecule;

• diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule.

• Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field:

Page 57: Molecular structure and bonding

Diatomic molecules

Page 58: Molecular structure and bonding

The energy level is the lower, the higher the EN of

the atom is !

Page 59: Molecular structure and bonding
Page 60: Molecular structure and bonding

Naming of MO’s: example O2 molecule

“g” = symmetric to C axis “u” = anti-symmetric

Page 61: Molecular structure and bonding

Diatomic molecules Consider the EN of each atom – the higher the EN, the lower is the energy of the orbitals ! The highest filled MO is called “HOMO”, the lowest unoccupied MO “LUMO” -> check example CO

http://firstyear.chem.usyd.edu.au/calculators/ mo_diagrams.shtml

Page 62: Molecular structure and bonding

Example CO

HOMO

LUMO

“lone pair” on C

Page 63: Molecular structure and bonding

Chemical Reactivity Important are the HOMO and LUMO (“frontier orbitals”)

http://www.meta-synthesis.com/webbook/12_lab/lab.html

Page 64: Molecular structure and bonding

Gro

up

Orb

itals

Page 65: Molecular structure and bonding

Construction of Group Orbitals – example H2O

Page 66: Molecular structure and bonding

Interaction 1: in-phase H orbitals

Page 67: Molecular structure and bonding

Interaction 2: out-of-phase H orbitals

Page 68: Molecular structure and bonding

Indicate different MO types: (bonding, non-bonding. anti-bonding)

Page 69: Molecular structure and bonding

Combination of 3 H orbitals to 3 group orbitals

BH3 molecule

Page 70: Molecular structure and bonding

Compare HOMO/LUMO to BH3 ! => what is an acid / base ?

Page 71: Molecular structure and bonding

Homework (3)

Number Molecule

1 CN

2 CN(-)

3 BC

4 BN

5 BO

6 BF

7 CF

8 NO

9 NO (+)

10 NO (-)

Number Molecule

11 NF

12 OF

13 CH4

14 BH3

15 SbF6

16 XeF2

17 XeF4

18 XeF6

http://firstyear.chem.usyd.edu.au/calculators/mo_diagrams.shtml