Chemistry 100 Chapter 9 Molecular Geometry and Bonding Theories

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 minimized 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

Four Electron Domains – Three Different Geometries Replacement of bonding domains (B)

with nonbonding domains (E)results in a different molecular geometry.

AB4 AB3E AB2E2

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 Diatomic Molecules

Hydrogen molecule a single bond between the

two H 1s orbitals a bond

Hydrogen Chloride a single bond from the

overlap of the Cl 3p orbital with the H 1s orbital

Chlorine molecule a single bond from the

overlap of the Cl 3p orbitals

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

Examine BH3 (a trigonal planar molecule)

sp Hybridisation

Examine BeF2 (a linear molecule). These sp hybrid orbitals have an

angle of 180 between them.

A Linear Molecule

The BeF2 molecule

Double Bonds

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

the bonding picture must be consistent with VSEPR theory.

Sigma () Bonds

Sigma bonds are characterized by Head-to-head overlap. Cylindrical symmetry of electron density

about the internuclear axis.

Pi () Bonds

Pi bonds are characterized by Side-to-side

overlap. Electron density

above and below the internuclear axis.

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 ) ] p type

The C2H4 Molecule

The Bond Angles in C2H4

Bond angles HCH = HCC 120. p 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 p bond

The Triple Bond in C2H2

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

the molecular plane.

Triple bond one bond and two p 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 ) ] p type [2pz

(C 1 ) – 2pz (C 2 ) ] p type

Bonding in H2O

Bonding Overlaps[sp3(O)–1s(H)] x 2

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) ] p 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 stype bonds).

5 sp3d orbitals ® these orbitals overlap with the 3p orbitals in Cl to form the 5 s 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 s bonds with the required VSEPR geometry ® octahedral.

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

Notes for Understanding Hybridization

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.

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

Delocalised Bonding

Valence bond theory – bonding electrons have been totally associated with

the two atoms that form the bond they are localized. What about the bonding situation in benzene,

the nitrate ion, the carbonate ion?

Bonding in Aromatic Molecules

Benzene C-C s bonds are formed from the sp2 hybrid orbitals. Unhybridized 2pz orbital on adjacent C atoms overlap

(bonds).

Bonding in the Benzene Molecule The p bonds extend over the whole

molecule the p electrons bonds are delocalized – they are

free to move around the benzene ring. Resonance structures – delocalization of the

-electrons.

The Nitrate Anion

Three resonance structures Alternating single

and double bonds Blend resonance

structures Delocalized

bond over anion backbone

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.

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

s1s

s*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.