Chemistry 100 Chapter 9

Molecular Geometry and Bonding Theories

Molecular Geometry

The three-dimensional arrangement of atoms in a molecule molecular geometry

Lewis structures can’t be used to predict geometry

Repulsion between electron pairs (both bonding and non-bonding) helps account for the molecular structure!

The VSEPR Model

Electrons are negatively charged, they want to occupy positions such that electron Electron interactions are minimised as much as possible

Valence Shell Electron-Pair Repulsion Model treat double and triple bonds as single domains resonance structure - apply VSEPR to any of them formal charges are usually omitted

Molecules With More Than One Central Atom

We simply apply VSEPR to each ‘central atom’ in the molecule.

• Carbon #1 – tetrahedral

• Carbon #2 – trigonal planar

Dipole Moments

The HF molecule has a bond dipole – a charge separation due to the electronegativity difference between F and H.

The shape of a molecule and the magnitude of the bond dipole(s) can give the molecule an overall degree of polarity dipole moment.

+H-F

Homonuclear diatomics no dipole moment (O2, F2, Cl2, etc)

Triatomic molecules (and greater). Must look at the net effect of all the bond dipoles.

In molecules like CCl4 (tetrahedral) BF3 (trigonal planar) all the individual bond dipoles cancel Þ no resultant dipole moment.

Bond Dipoles in Molecules

More Bond Dipoles

Valence Bond Theory and Hybridisation

Valence bond theory description of the covalent bonding and

structure in molecules. Electrons in a molecule occupy the

atomic orbitals of individual atoms. The covalent bond results from the

overlap of the atomic orbitals on the individual atoms

The Bonding in H2

Hydrogen molecule a single bond indicating the overlap of the

1s orbitals on the individual atoms cylindrical symmetry with respect to the

line joining the atomic centres, i.e., a bond

H H

Overlap Region

1s (H1) – 1s(H2) bond

The Bonding in H2

H H

The Cl2 Molecule

In the chlorine molecule, we observe a single bond indicating the overlap of the 3p orbitals on the individual atoms.

Cl Cl

Bonding description 3pz (Cl 1) – 3pz (Cl 2)

Is This a Bond?

Cl Cl

Hybrid Atomic Orbitals

Look at the bonding picture in methane (CH4).

• Bonding and geometry in polyatomic molecules may be explained in terms of the formation of hybrid atomic orbitals

• Bonds overlap of the hybrid atomic orbitals on central atoms with appropriate half-filled atomic orbital on the terminal atoms.

The CH4 Molecule

The Formation of the sp3 Hybrids

Mix 3 “pure” p orbitals and a “pure” s orbital form an sp3

“hybrid” orbital. Rationalize the

bonding around the C central atom.

sp2 Hybridisation

What if we try to rationalise the bonding picture in the BH3 (a trigonal planar molecule)?

We mix 2 “pure” p orbitals and a “pure” s orbital to form “hybrid” or mixed sp2 orbitals.

These three sp2 hybrid orbitals lie in the same plane with an angle of 120 between them.

A Trigonal Planar Molecule

Overlap regions

B

H

H H

Overlap region

sp Hybridisation

What if we try to rationalize the bonding picture in the BeH2 species (a linear molecule)?

We mix a single “pure” p orbital and a “pure” s orbital to form two “hybrid” or mixed sp orbitals

These sp hybrid orbitals have an angle of 180 between them.

A Linear Molecule

The BeH2 molecule

BeH H

Overlap Regions

Double Bonds

Look at ethene C2H4. Each central atom is an AB3 system,

the bonding picture must be consistent with VSEPR theory.

The Bond

Additional feature an unhybridized p

orbital on adjacent carbon atoms.

Overlap the two parallel 2pz orbitals (a -bond is formed).

Bond overlaps in C2H4

There are three different types of bonds

[sp2 (C ) – 1s (H) ] x 4 type

[sp2 (C 1 ) – sp2 (C 2 ) ] type

[2pz (C 1 ) – 2pz

(C 2 ) ] type

The C2H4 Molecule

The Bond Angles in C2H4

Bond angles HCH = HCC 120. bond is perpendicular

to the plane containing the molecule.

Double bonds – Rationalize by assuming

sp2 hybridization exists on the central atoms!

Any double bond one bond and a bond

The Triple Bond in C2H2

Bond angles HCH = HCC = 180. bonds are perpendicular to

the molecular plane. Triple bond one bond and

two bonds

Triple bond rationalized by assuming sp hybridization exists on

the central atoms!

Bond Overlaps in C2H2

There are again three different types of bonds[sp (C ) – 1s (H) ] x 2 type [sp (C 1 ) – sp (C 2 ) ] type

[2py (C 1 ) – 2py

(C 2 ) ] type [2pz

(C 1 ) – 2pz (C 2 ) ] type

Bond Overlaps in H2CO

There are again three different types of bonds[sp (C) – 1s (H) ] x 2 type [sp2 (C) – sp2 (O) ] type [2p (C) – 2p (O) ] type

Key Connection – VSEPR and Valence Bond Theory!!

sp3d Hybridisation

How can we use the hybridisation concept to explain the bonding picture PCl5.

There are five bonds between P and Cl (all type bonds).

5 sp3d orbitals these orbitals overlap with the 3p orbitals in Cl to form the 5 bonds with the required VSEPR geometry trigonal bipyramid.

Bond overlaps[sp3d (P ) – 3pz (Cl) ] x 5 type

sp3d2 Hybridisation

Look at the SF6 molecule. 6 sp3d2 orbitals these orbitals

overlap with the 2pz orbitals in F to form the 6 bonds with the required VSEPR geometry octahedral.

Bond overlaps [sp3d2 (S ) – 2pz (F) ] x 6 type

Notes for Understanding Hybridisation Applied to atoms in molecules only Number hybrid orbitals = number of atomic

orbitals used to make them Hybrid orbitals have different energies and

shapes from the atomic orbitals from which they were made.

Hybridisation requires energy for the promotion of the electron and the mixing of the orbitals energy is offset by bond formation.

Delocalised Bonding

In almost all the cases where we described the bonding n the molecule, the bonding electrons have been totally associated with the two atoms that form the bond they are localised.

What about the bonding situation in benzene, the nitrate ion, the carbonate ion?

In benzene, the C-C bonds are formed from the sp2 hybrid orbitals. The unhybridised 2pz orbital on C overlaps with another 2pz orbital on the adjacent C atom.

Three bonds are formed. These bonds extend over the whole molecule (i.e. the bonds are delocalised).

The electrons are free to move around the benzene ring.

Any species where we had several resonance structures, we would have delocalisation of the -electrons.

Delocalised Electrons in Molecules

Molecular Orbital (M.O.) Theory

Valence bond and the concept of the hybridisation of atomic orbitals does not account for a number of fundamental observations of chemistry.

To reconcile these and other differences, we turn to molecular orbital theory (MO theory).

MO theory – covalent bonding is described in terms of molecular orbitals the combination of atomic orbitals that results in

an orbital associated with the whole molecule.

Recall the wave properties of electrons.

constructive interference the two e- waves interact favourably; loosely analogous to a build-up of e- density between the two atomic centres. destructive interference unfavourable interaction of e- waves; analogous to the decrease of e- density between two atomic centres.

Constructive and Destructive Interference

+

+

Constructive

Destructive

ybonding = C1 ls (H 1) + C2 ls (H 2) yanti = C1 ls (H 1) - C2 ls (H 2)

Bonding Orbital a centro-symmetric orbital (i.e. symmetric about the line of symmetry of the bonding atoms).

Bonding M’s have lower energy and greater stability than the AO’s from which it was formed.

Electron density is concentrated in the region immediately between the bonding nuclei.

Anti-bonding orbital a node (0 electron density) between the two nuclei.

In an anti-bonding MO, we have higher energy and less stability than the atomic orbitals from which it was formed.

As with valance bond theory (hybridisation)2 AO’s 2 MO’s

Bonding and Anti-Bonding M.O.’s from 1s atomic Orbitals

Energy

1s

1s

1s

*1s

The MO’s in the H2 Atom

The situation for two 2s orbitals is the same! The situation for two 3s orbital is the same.

Let’s look at the following series of moleculesH2, He2

+, He2

bond order = ½ {bonding - anti-bonding e-‘s}. Higher bond order º greater bond stability.