Molecular Structure and Bonding

Dr.Christoph Phayao University March 2015

Part 1

What is a chemical bond ?

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

Covalent Bond

Two non-metals share (valence) electrons:

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

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

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

Bond Polarity

Polar Bonds

Uneven sharing of electrons due to differences in Electronegativity

The “pull” an atom has for electrons

Electronegativity Trends

Common Electronegativites

Highest value, set to 4

“Calculate” EN The first idea of EN came from L.Pauling – he estimated EN from bond energies

Allred-Rochow EN

Example: Flourine (r = 72 pm) Carbon (r = 77 pm) Calculate the AR electronegativity

Shielding of 2p electrons: Flourine: S = 6 * 0.35 + 2 * 0.85 = 3.8 => Z* = 9 – 3.8 = 5.2 EN = 3590 * 5.2/(72 2) + 0.744 = 4.35 Carbon: S = 3 * 0.35 + 2 * 0.85 = 2.75 => Z* = 6 – 2.75 = 3.25 EN = 3590 * 3.25 / (77 2) ) + 0.744 = 2.71

Calculate Dipole moments

Polar Molecules

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

Polar Bonds and Geometry

Which bond type ?

(exception: Transition metals !)

Electron counting

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

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 !

Part 2: Valence Bond Theory (VB)

“Valence Electrons are located in bonds and lone pairs”

Sigma bonds

Pi Bonds

“Resonance”

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

Important exception: Carbon Monoxide !

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

***** Break *****

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

Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.

atomic

orbitals

hybrid

orbitals

orbital box diagrams

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

orbital box diagrams with orbital contours

Figure 11.3 The sp2 hybrid orbitals in BF3.

sp2 and sp3

Hybrid Orbitals

Figure 11.4 The sp3 hybrid orbitals in CH4.

Figure 11.5 The sp3 hybrid orbitals in NH3.

Figure 11.5 continued The sp3 hybrid orbitals in H2O.

Including d-orbitals

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

Figure 11.6 The sp3d hybrid orbitals in PCl5.

Figure 11.7 The sp3d2 hybrid orbitals in SF6.

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

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

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 !

Acid or Base ?

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

***** Break *****

VSEPR

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

Practise: http://www.youtube.com/watch?v=xwgid9YuH58

Rules to remember

Order of repulsions:

(1) Lone Pair – Lone Pair

(2) Lone Pair – Bond

(3) Bond - Bond

Examples

XeF2

The lone pair needs most space, so they are in equatorial position (bond angle 120 deg) ClF3 lone pairs again in eq. position so have max. distance

Estimate the structures of:

SF4

BrF5

IF5 2-

AsF5

cis and trans

mer and fac

eq and ax

MO Theory

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).

Amplitudes of wave

functions added

Figure 11.14

An analogy between light waves and atomic wave functions.

Amplitudes of

wave functions

subtracted.

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.

Prentice Hall © 2003 Chapter 9

Molecular Orbitals

The Hydrogen Molecule

Prentice Hall © 2003 Chapter 9

Figure 11.15 The MO diagram for H2.

Energ

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.

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:

Diatomic molecules

Please try to draw the AO’s

and MO’s

Naming of MO’s: example O2 molecule

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

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

Example CO

HOMO

LUMO

“lone pair” on C

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

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

Gro

up

Orb

itals

Construction of Group Orbitals – example H2O

Interaction 1: in-phase H orbitals

Interaction 2: out-of-phase H orbitals

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

Combination of 3 H orbitals to 3 group orbitals

BH3 molecule

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

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