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. H 2 O, NH 3, CH 4, O 2, N 2, SCl 2, CCl 4, PCl 3, NO +, NH 4 +, CO, CO 2

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 AB n 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 spell the names!

How to predict geometry of a molecule? 1)Draw the Lewis structure of the molecule or ion, and count the number of electron domains around the central atom. 2)Determine the electron domain arrangement by arranging the electron domains about the central atom so that the repulsions among them are minimized. 3) Use the arrangement of the bonded atoms to determine the molecular geometry. CO 2

BF 3, NO 3 −, H 2 CO

Electron domains for multiple bonds exert a greater repulsion force on adjacent electron domains than do electron domains for single bonds. lone pair-lone pair > lone pair-bonding pair > bonding pair- bonding pair

SO 2 119° Electron domain arrangement is not necessarily the same as the molecular structure. Bent or V-shaped

CH 4

NH 3

H2OH2O

PCl 5

SF 4 To minimize repulsion, electron lone pairs are always placed in equatorial positions for trigonal bipyramidal geometry.

BrF 3

XeF 2

SF 6

BrF 5

XeF 4

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 — F dipole 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

v1v1 v2v2 v = v 1 + v 2 v 3 = v 2

SO 3

120° = = Net dipole moment = 0, nonpolar molecule SO 3

109.5° CCl 4 = = Net dipole moment = 0, nonpolar molecule

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

NH 3

How are electrons shared in covalent bonds? Valence Bond Theory Molecular Orbital Theory

Valence Bond Theory: Orbital Overlap as a Chemical Bond

CH 4

4 electron domains sp 3 hybridization tetrahedral arrangement

NH 3

sp 3 on O H2OH2O

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 sp 2 hybridization trigonal planar arrangement

H−C ≡ C−H Acetylene Linear molecule

N2N2 :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 OC ≡ N: :: ¨ ¨

PCl 5

The Orbitals Used to Form the Bonds in PCl 5

5 electron domains dsp 3 hybridization trigonal bipyramidal arrangement

SF 6

An Octahedral Set of d 2 sp 3 Orbitals on Sulfur Atom

6 electron domains d 2 sp 3 hybridization octahedral arrangement

paramagnetismdiamagnetism paired electronsunpaired electrons

Liquid O 2 is paramagnetic

How are electrons shared in covalent bonds? Valence Bond Theory Molecular Orbital Theory

Atoms → atomic orbitals Molecules → molecular orbitals Molecular Orbital (MO) ≈ Linear Combination of Atomic Orbitals (LCAO)

H2H2

Electron configuration (σ 1s ) 2 Pauli principle and Hund’s rule apply

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 (σ 1s ) 2 bond order = 1

(σ 1s ) 2 (σ 1s * ) 1 bond order = 0.5

(σ 1s ) 2 (σ 1s * ) 2 bond order = 0

(σ 1s ) 2 (σ 1s * ) 1 bond order = 0.5

Li 2

(σ 2s ) 2 bond order = 1

(σ 2s ) 2 (σ 2s * ) 2 bond order = 0

O 2, F 2, Ne 2 : (σ 2s ) (σ 2s * ) (σ 2p ) (π 2p ) (π 2p * ) (σ 2p * ) B 2, C 2, N 2 : (σ 2s ) (σ 2s * ) (π 2p ) (σ 2p ) (π 2p * ) (σ 2p * ) 1)Electron configuration. 2)Bond order → stable molecule/ion? 3)Paramagnetic or Diamagnetic? O 2, F 2, N 2, N 2 −, N 2 +

Homonuclear diatomic molecules/ions Heteronuclear diatomic molecules/ions

Problems Chapter ,33,35,39,42,47,51,53,59,61,63,69, 71,85,86