Chapter 10
 Molecular Geometries and Bonding
Theories
– Lewis structures do not indicate the molecular architecture – the shape of the molecule.
– The shape and structure of a molecule determines much of its physical and chemical characteristics.

VSEPR Theory
 Valence-shell
Electron Pair
Repulsion
– Electron pairs
(domains or regions) repel each other completely.
– Balloon model.

Electron Regions

1.
2.
3.
The number of electron regions around the central atom are counted as:
Each single bond counts as a region.
Each lone pair counts as a region.
A multiple bond counts as a single region.

Electron Regions
 How many?

Electron Pair Geometry (EPG)


Can be from two to six regions.
Thus, only five EPG’s are possible.
 Two regions produces a linear EPG.

Electron Pair Geometry
 Three regions produces a trigonal planar geometry.
 Planar = 2D.
 Ex) BF
3

Electron Pair Geometry
•
Four regions becomes a three dimensional structure based on the tetrahedron.
•
Formally called tetrahedral with bond angles of 109.5
o

Electron Pair Geometry
 Tetrahedral is very common and symmetrical.
 An example is CF
4

Electron Pair Geometry
 Five regions produces a trigonal bipyramidal geometry with two sets of bond angles.

Electron Pair Geometry
 An example is PCl
5

Electron Pair Geometry
 Six regions produces an octahedral geometry.

Electron Pair Geometry
 An example is SF
6

Molecular Geometry (MG)
 This is based on the shape of the electron pairs.
 When a molecule has no lone pairs, the
EPG = MG.
 If the molecules DOES have one or more lone pairs, then the shape of the atoms is determined based off of the
EPG.

Molecular Geometry

Molecular Geometry

Molecular Geometry

Examples








Bent (120), SO
2
Trigonal pyramidal, NH
3
Bent (109.5), H
2
O
Seesaw, SF
4
T-shaped, ClF
3
Linear, I
3
-
Square pyramidal, BrF
5
Square planar, XeF
4

Sketching the Molecules
 Simple = Ball and Stick figures
 Representing the 3D shapes:
– Put as many of the molecules in the same plane as possible including the central atom. Use straight lines for bonds connected to atoms in plane.
– For atoms in front of the plane, use a solid wedge.
– For atoms behind the plane, use a hashed wedge.

3D Representations

Lone Pairs
 A non-bonding pair will always take up more space.
 This compresses the normal bond angles.

Lone Pairs

Lone Pairs
 This also explains the MG’s for the trigonal bipyramidal family.

Shapes of Larger Molecules
 A molecule like acetic acid has three central atoms.

 A molecule can contain very polar bonds, but can be non-polar.
 An example is
CO
2
.

Polarity
 On the other hand, sometimes polar bonds DO make a molecule polar.
 An example of a polar molecule is
H
2
O.

 A molecule with a symmetrical distribution of polar bonds will be nonpolar.
 A molecule with an un-symmetrical distribution of polar bonds will be polar.
– presence of lone pairs
– different external atoms

Polarity

Polarity
 Polar molecules are attracted to other polar molecules
 Because water is a polar molecule, other polar molecules dissolve well in water
– and ionic compounds as well
 Non-polar molecules do
NOT dissolve in water.

Valence Bond Theory
 How can we explain the formation of the bonds in a molecular compound?
 A bond occurs when a valence orbital on one atom overlaps with a valence orbital of another atom.

 The H
2 molecule – a closer look.
nuclear repulsion no interaction minimum energy

Valence Bond Theory
 Three (or more) atom molecules cannot be explained by simple overlap of orbitals.
 Fact: a bond generally forms between two half-filled orbitals.
 Fact: an s-type orbital is spherical, so it could form a bond in any direction.
 Fact: the three p-type orbitals are at 90 degree angles to each other.

 CH
4
– has an EPG and
MG of tetrahedral with bond angles of 109.5
o .
 Valence diagram for C and H before any bonding is:

 Solution: promote the paired electron from the s orbital to the empty p orbital.
 Solution: mix the one s and three p orbitals together to get a new set of four orbitals all equal in energy. This is called _____________________.

Valence Bond Theory
 Each hybrid orbital has some s and some p characteristics.
 Thus, they look different!

Types of Hybrids
 Determined from the EPG.
EPG
Linear
Atomic orbitals s+p = sp
Hybrid diagram
Trigonal planar s+p+p = sp 2
Tetrahedral s+p+p+p = sp 3
Examples
BeF
2
BF
3
CH
4

Types of Hybrids
 Atoms in the third period and beyond have empty d orbitals that can potentially be used for hybridization.
 PCl
5
– requires five bonds, so need a set of five orbitals.
 Once again, must first promote the s electron to an empty d orbital.

Types of Hybrids
EPG Atomic orbitals
Trigonal bipyramidal s+p+p+p+d
= sp 3 d
Hybrid diagram
Examples
PCl
5
Octahedral s+p+p+p+d+d = sp 3 d 2
SF
6

Molecules with Lone Pairs
 Ex) NH
3
 Ex) H
2
O
 Ex) BrF
3

Multiple Bonds


Two types of bonds are possible.
1. Sigma ( s
) bonds have a cylindrical shape of electron density along the central axis between the two nuclei.
s bond

Multiple Bonds
 2. Pi ( p
) bonds have an electron density above and below the central axis.
– Are formed by the overlap of two parallel half-filled p-type orbitals.

Multiple Bonds
 The majority of bonds are sigma bonds.
 When a double bond is present, the first bond is a sigma and the second is a pi.

Pi bonds

Multiple Bonds
 For any pi bonds, you MUST use an unhybridized half-full p-type orbital.
 Ex) C
2
H
4
 Ex) CO
2

Multiple Bonds

Pi Bond Significance
 Sigma bonds have free-rotation about the central axis.
 Ex) C
2
H
4
Cl
2
 Pi bonds have NO free-rotation due to the fact that they must overlap above and below the central axis.
 Ex) C
2
H
2
Cl
2

Pi Bond Significance

Isomers
 When two compounds share the exact same formula but are different either structurally or spatially, then they are said to be isomers.
 Structural isomers
– C
5
H
12
– C
2
H
6
O

Isomers
 Geometric isomers are different spatially.
 This can occur for our carbon-carbon double bond.
X Y
C=C
Y
C=C
Y
Y X X X
Trans Cis

Isomers
 The last molecule in your packet has three possible structures. One is structural and two are geometric isomers.
 One other geometry can have cis/trans isomerism – is it tetrahedral or square planar?
 Ex) CH
2
Cl
2 or Pt(NH
3
)
2
Cl
2
?

Limitations of V.B. Theory
 Valence Bond Theory does not adequately explain molecules with resonance structures nor some other observed properties.
 Ex) O
2 or molecular oxygen is paramagnetic (unpaired electrons).
Lewis structure for O
2

Molecular Orbital Theory
 A more sophisticated and complex model of bonding.
 Atomic orbitals from each atom contribute to new MO’s.
 Like atomic orbitals, each MO can hold up to two electrons.
 A MO, though, is spread out over the entire molecule.

MO Theory
 For each atomic orbital contributed we get one MO.
 Half of the MO’s become bonding and the other half become anti-bonding.
– Waveforms add either constructively or destructively like light!
 For the n=1 period, each atom contributes a 1s atomic orbital.

MO Theory

MO Theory

MO Theory
 The Bond Order in MO theory is found by: BO = ½ (Bonding e - Anti-bonding e ).
 Any bond order = 0 implies that the molecule is not possible.
 Odd number of electrons will produce halfinteger BO’s.

MO Theory
 Period 2 elements have both the 2s and
2p atomic orbitals to contribute towards
MO’s.
 Thus, two atoms from period 2 will have how many atomic orbitals total?
 How many MO bonding orbitals will be produced? Anti-bonding?

MO Theory
 The two 2s orbitals overlap just like the
 two 1s orbitals did in period 1.
This produces the s
2s and s
2s
* MO’s.
 The six 2p orbitals overlap differently.
– Two will overlap end-on-end and produce a s
2p type MO.
– Four will overlap sideways and produce two p MO’s that are equal in energy.

MO Theory
 The energy level diagram produced for all of these new
MO’s is:

MO Theory
 Diagram assumes that no 2s-2p orbitals interactions occur.
 For B s
2p
2
, C
2 and p
2p
, and N
2 the interactions cause the order to trade places on the diagram.
 Since these are filled for O
2
, F
2
, and Ne
2 the diagram can be written with those two always reversed to simplify.

MO Theory

MO Theory


It is possible for some molecules and ions to be paramagnetic – one or more unpaired electrons.
Most will be diamagnetic – all paired electrons.

n=2 Diatomic Molecules

n=2 Diatomic Molecules

Molecular Oxygen
 According to MO theory, it has a BO = 2 and it is paramagnetic!
 As liquid O
2 is poured between the poles of a magnetic, it will have a strong attraction.
 Clip.

Heteronuclear Diatomics
 For period 2, we can mix and match other elements and apply the same MO diagram.
 Simply add up the total valence electrons that each contribute and place in the diagram.
 For ions, a positive charge means we would decrease by one electron and a negative charge means we would increase by one electron.

Heteronuclear Diatomics
 More electronegative element has lower energy orbitals.
 Can produce bond orders with ½ values.
 B.O. = ___________

Polyatomic Molecules
 When many atoms are combined together, the atomic orbitals of all the atoms are combined to make a set of molecular orbitals, which are delocalized over the entire molecule
 Gives results that better match real molecule properties than either Lewis or valence bond theories
 This is why resonance structures cannot be explained by valence bond theory.

Ozone, O
3
 The structure of O
3 includes two resonances.
M.O. showing delocalized pi bonds