Chapter 9Molecular Shapes

-shape of molecule is based on bond angles

Valence Shell Electron Pair Repulsion (VSEPR)

-based on the idea that electron groups repel one another

-for molecules having one central atom, shape is based on:

e- domain- region where e- are likely to be found

one e- domain = a lone pair, single, double or triple bond

-must have the structural formula drawn

-use the following formula

ABnEn

A= central atom

B= # of bonding domains on central atom

E= # of lone pairs of e- on central atom

e- geometry- shows arrangement of e- domains

molecular geometry- arrangement of the atoms

Molecular Shape and Polarity

-to determine if an entire molecule is polar you must look at the bonds and the shape

dipole moment- a measure of the separation of + and – charges in a molecule

bond dipole- dipole moment due only to the two atoms in that bond

-bond dipoles and dipole moments are vector quantities

-have both magnitude and direction and must look at both when determining polarity

-vectors will be added

To determine if a molecule is polar:

-determine the shape

-determine whether the molecule contains polar bonds

-determine whether the polar bonds add together to form a net dipole moment

-if they sum to zero then the molecule is

nonpolar

-if not then the molecule is polar

-if the molecule has lone pairs, it is polar

Tro, Chemistry: A Molecular Approach

7

Examples:a) BrCℓ

polarb) SO2

polar *has lone pairc) SF6

nonpolard) NF3

polar *lone paire) BCℓ3

nonpolar

Valence Bond Theory

-valence e- of atoms in a molecule reside in atomic orbitals which can be s, p, d, f or some hybrid of these

-a chemical bond results from the overlap of two half-filled orbitals, or less commonly from the overlap of a completely filled orbital with an empty orbital

-the shape of the molecule is determined by the geometry of the overlapping orbitals

Ex: H2S

-overlap of half-filled orbitals does not explain bonding in all types of molecules

-for example: CH4

-b/c there are only two half-filled orbitals, we would assume that C bonds with only two H

-we know it bonds with 4 H

-we must consider hybrid orbitals here

hybridization- standard atomic orbitals are combined to form hybrid orbitals that correspond more closely to the actual distribution of e- in chemical bonds

General Statements Regarding Hybridization

-# of standard atomic orbitals added together = # of hybrid orbitals formed

-the combo of orbitals determines the shapes and energies of orbitals formed

-the specific type of hybridization that forms for a molecule is the one that yields the lowest overall energy for the molecule

-a single bond contains a sigma bond (σ)- head on overlap

-a double bond is made up of one sigma and one pi bond (π)- sideways overlap

-a triple bond is made up of one sigma and two pi bonds

sp3 hybridization

-occurs when there are 4 e- groups (tetrahedral)

-for CH4, one s orbital and three p orbitals hybridize to form four sp3 orbitals of equal energy

-C has four half-filled orbitals to overlap with four half-filled H 1s orbitals

*all are single bonds sigma bonds (σ)

sp2 hybridization and double bonds

-hybridization of one s orbital and two p orbitals results in three sp2 hybrids and one leftover unhybridized p orbital

-occurs when you have 3 e- groups (tri. planar)

Example: H2CO

-because C is the central atom it undergoes hybridization

*made up of two sigma bonds from the C-H single bonds, and one sigma and one pi bond (π) from the double bond of C=O

**Try H2C2Cℓ2

sp hybridization and triple bonds

-one s orbital and one p orbital hybridize to form two sp orbitals of equal energy

-two p orbitals remain unhybridized

-occurs when there are two e- groups (linear)

Example: C2H2

-has a triple bond

Practice:

Tell the type of hybridization of the C atom(s), what types of bonds, and what kinds of orbitals for each.

1) H3C2OH 2) HCN 3) C2H4

1) 1st C= sp3, 4 σ bonds, 1s and 3p orbitals

2nd C= sp2, 3 σ and 1 π bond, 1s, 2p, and 1 unhybridized p orbital

2) sp, 2 σ bonds, 2 π bond, 1s, 1p and 2 unhybridized p orbitals

3) **both interior C the same= sp2, 3 σ bonds, 1 π bond, 1s, 2p, and 1 unhybridized p orbital

-if trigonal bypyramidal, hybridization is sp3d

-if octahedral, hybridization is sp3d2

localization of e-: σ and π e- are associated only with the two atoms that form the bond

delocalization of e-: e- in π bonds extend over more than just the two atoms

*occurs when a molecule has at least two resonance structures

Molecular Orbital Theory

-orbitals are treated as overlapping the entire molecule, not as “belonging” to individual atoms

-when atomic orbitals are combined to form molecular orbitals, the total number of orbitals does not change

-any time you bond two atomic orbitals you get two molecular orbitals:

bonding orbital (σ or π ) - lower in energy

anti-bonding (σ* or π* ) - higher in energy

*b/c lower energy bonding fills first

*each orbital can hold two e-

*will show if a single, double or triple bond is formed by calculating the bond order

bond order = # e- in bonding - # e- in antibonding

2

-if bond order is zero or less, no bond will form

-σ are one box b/c of head on overlap

-π are two boxes b/c of sideways overlap

s orbitals = σ and σ* for the bonding and antibonding

p orbitals = π, σ, π* and σ* for bonding and antibonding

paramagnetism- having unpaired e-

diamagnetism- having no unpaired e-