Chapter 10 Chemical Bonding II. Lewis Structure Molecular Structure Structure determines chemical properties.

Chapter 10

Chemical Bonding II

Lewis Structure Molecular Structure

Structure determines chemical properties

Electron domain/group: area where electrons appear in Lewis structures.

It can be electron lone pairs, single bonds,

double bonds, triple bonds, or single electrons.

H2O, NH3, CH4, O2, N2, SCl2, CCl4, PCl3, NO+, NH4

+, CO, CO2

Valence Shell Electron Pair Repulsion (VSEPR) model

The lowest energy arrangement of a given number of electron

domains is the one that minimizes the repulsions among them.

The shape of ABn molecules or ions depend on the number of

electron domains surrounding the central A atom.

number of electron domains: 2 to 6

Know how to spellthe names!

How to predict geometry of a molecule?

1) Draw the Lewis structure of the molecule or ion, and countthe number of electron domains around the central atom.

2) Determine the electron domain arrangement by arranging theelectron domains about the central atom so that the repulsionsamong them are minimized.

3) Use the arrangement of the bonded atoms to determine the molecular geometry.

CO2

BF3, NO3−, H2CO

Electron domains for multiple bonds exert a greaterrepulsion force on adjacent electron domains than do electron domains for single bonds.

lone pair-lone pair > lone pair-bonding pair > bonding pair- bonding pair

SO2

119°

Electron domain arrangement is not necessarily the same as the molecular structure.

Bent or V-shaped

CH4

NH3

H2O

PCl5

SF4

To minimize repulsion, electron lone pairs are always placed in equatorial positions for trigonal bipyramidal geometry.

BrF3

XeF2

SF6

BrF5

XeF4

Polarity of a molecule

How to quantify the polarity of a bond?

Dipole moment

+ −Dipole

Dipole has a magnitude and a direction — vector

Magnitude (length) of a dipole — dipole moment

μ = qr

q — charge, r — distance between + and − charge

H — Fdipole

dipole moment of a bond ≠ 0 ↔ polar bond

dipole moment of a bond = 0 ↔ nonpolar bond

The Pauling Electronegativity Values

Polarity of a molecule

dipole moment of a molecule ≠ 0 ↔ polar molecule

dipole moment of a molecule = 0 ↔ nonpolar molecule

dipole of a molecule = sum of all the bond dipoles

v1

v2

v = v1 + v2

v3 = v2

SO3

120°

=

=

Net dipole moment = 0, nonpolar molecule

SO3

109.5°

CCl4

=

=Net dipole moment = 0, nonpolar molecule

Polarity of a molecule depends on the polarity of its bonds AND the geometry of the molecule.

NH3

How are electrons shared in covalent bonds?

Valence Bond Theory

Molecular Orbital Theory

Valence Bond Theory:

Orbital Overlap as a Chemical Bond

CH4

CH4

4 electron domains

sp3 hybridization

tetrahedral arrangement

NH3

sp3 on OH2O

C = C

Ethylene

H

H

H

H

~120°

All atoms are in the same plane

C = C

Ethylene

H

H

H

H

~120°

All atoms are in the same plane

The σ Bonds in Ethylene

C = C

Ethylene

H

H

H

H

~120°

All atoms are in the same plane

A Carbon-Carbon Double Bond Consists of a σ and a π Bond

3 electron domains

sp2 hybridization

trigonal planar arrangement

H−C ≡ C−H

Acetylene

Linear molecule

N2 :N≡N:

2 electron domains

sp hybridization

Linear arrangement

Single bond: σ

double bond: one σ, one π

triple bond: one σ, two π

C

H

H

CH

=O

O C ≡ N:

::

¨¨

PCl5

The Orbitals Used to Form the Bonds in PCl5

5 electron domains

dsp3 hybridization

trigonal bipyramidal arrangement

SF6

An Octahedral Set of d2sp3 Orbitals on Sulfur Atom

6 electron domains

d2sp3 hybridization

octahedral arrangement

paramagnetismdiamagnetism

paired electrons unpaired electrons

Liquid O2 is paramagnetic

How are electrons shared in covalent bonds?

Valence Bond Theory

Molecular Orbital Theory

EH

Atoms → atomic orbitals

Molecules → molecular orbitals

Molecular Orbital (MO)≈ Linear Combination of Atomic Orbitals (LCAO)

H2

Electron configuration

(σ1s)2

Pauli principleand Hund’s ruleapply

Bond order = ½ (number of bonding electrons − number of antibonding electrons)

If bond order > 0, the molecule is stable

If bond order = 0, the molecule is not stable

bond order = 1 → single bond

bond order = 2 → double bond

bond order = 3 → triple bond

Electron configuration

(σ1s)2 bond order = 1

(σ1s)2(σ1s*)1 bond order = 0.5

(σ1s)2(σ1s*)2 bond order = 0

(σ1s)2(σ1s*)1 bond order = 0.5

Li2

(σ2s)2 bond order = 1

(σ2s)2(σ2s*)2 bond order = 0

O2, F2, Ne2: (σ2s) (σ2s*) (σ2p) (π2p) (π2p

*) (σ2p*)

B2, C2, N2: (σ2s) (σ2s*) (π2p) (σ2p) (π2p

*) (σ2p*)

1)Electron configuration.

2)Bond order → stable molecule/ion?

3)Paramagnetic or Diamagnetic?

O2, F2, N2, N2−, N2

+

Homonuclear diatomic molecules/ions

Heteronuclear diatomic molecules/ions

Problems Chapter 10

1-7 ,33,35,39,42,47,51,53,59,61,63,69,

71,85,86