1. Chapter 10 Chemical Bonding II

2. Structure Determines Properties!
properties of molecular substances depend on the structure of the molecule
the structure includes many factors, including:
the skeletal arrangement of the atoms
the kind of bonding between the atoms
ionic, polar covalent, or covalent
the shape of the molecule
bonding theory should allow you to predict the shapes of molecules

3. Molecular Geometry Molecules are 3-dimensional objects
We often describe the shape of a molecule with terms that relate to geometric figures
These geometric figures have characteristic “corners” that indicate the positions of the surrounding atoms around a central atom in the center of the geometric figure
The geometric figures also have characteristic angles that we call bond angles

4. Using Lewis Theory to Predict Molecular Shapes
Lewis theory predicts there are regions of electrons in an atom based on placing shared pairs of valence electrons between bonding nuclei and unshared valence electrons located on single nuclei
this idea can then be extended to predict the shapes of molecules by realizing these regions are all negatively charged and should repel

5. VSEPR Theory
electron groups around the central atom will be most stable when they are as far apart as possible – we call this valence shell electron pair repulsion theory
since electrons are negatively charged, they should be most stable when they are separated as much as possible
the resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule

6. Electron Groups
the Lewis structure predicts the arrangement of valence electrons around the central atom(s)
each lone pair of electrons constitutes one electron group on a central atom
each bond constitutes one electron group on a central atom
regardless of whether it is single, double, or triple
there are 3 electron groups on N
1 lone pair
1 single bond
1 double bond

7. Molecular Geometries
there are 5 basic arrangements of electron groups around a central atom
based on a maximum of 6 bonding electron groups
though there may be more than 6 on very large atoms, it is very rare
each of these 5 basic arrangements results in 5 different basic molecular shapes
in order for the molecular shape and bond angles to be a “perfect” geometric figure, all the electron groups must be bonds and all the bonds must be equivalent
for molecules that exhibit resonance, it doesn’t matter which resonance form you use – the molecular geometry will be the same

8. Parent electronic structure

9. Examples
How many electron groups (charge clouds) are around the central atom in the following?
SO2 NH4+ PCl5

10. Trigonal Bipyramidal Geometry
when there are 5 electron groups around the central atom, they will occupy positions in the shape of a two tetrahedral that are base-to-base with the central atom in the center of the shared bases
this results in the molecule taking a trigonal bipyramidal geometry
the positions above and below the central atom are called the axial positions
the positions in the same base plane as the central atom are called the equatorial positions
the bond angle between equatorial positions is 120°
the bond angle between axial and equatorial positions is 90°

11. Octahedral Geometry
when there are 6 electron groups around the central atom, they will occupy positions in the shape of two square-base pyramids that are base-to-base with the central atom in the center of the shared bases
this results in the molecule taking an octahedral geometry
it is called octahedral because the geometric figure has 8 sides
all positions are equivalent
the bond angle is 90°

12. The Effect of Lone Pairs
lone pair groups “occupy more space” on the central atom
because their electron density is exclusively on the central atom rather than shared like bonding electron groups
relative sizes of repulsive force interactions is:
Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair
this effects the bond angles, making them smaller than expected
The bonding electrons are shared by two atoms, so some of the negative charge is removed from the central atom.
The nonbonding electrons are localized on the central atom, so area of negative charge takes more space

13. Derivative of Trigonal Geometry
when there are 3 electron groups around the central atom, and 1 of them is a lone pair, the resulting shape of the molecule is called a trigonal planar - bent shape
the bond angle is < 120°

14. Bond Angle Distortion from Lone Pairs

15. Replacing Atoms with Lone Pairs in the Trigonal Bipyramid System

16. T-Shape
Linear Shape

17. Predicting the Shapes Around Central Atoms
Total # of e- groups on central atom
“Parent” electronic geometry
# Bonded
atoms
# Lone
pairs
Idealized molecular shape
Idealized bond angles 2
Linear
180o 3
Trigonal Planar
120 o 1
Bent 4
Tetrahedral
109.5 o
Trigonal Pyramidal 5
Trigonal Bipyramidal
90 o, 120 o, 180 o
Seesaw
T-shaped
90 o, 180 o
180 o 6
Octahedral
Square Pyramidal
Square Planar

18. Real bond angles vs. Idealized bond angles
VSEPR predicts the idealized bond angle(s) by assuming that all electron groups take up the same amount of space. Since lone pairs are attracted to only one nucleus, they expand into space further than bonding pairs, which are attracted to two nuclei. As a result, real molecules that has lone pairs on the central atom often have bond angles that are slightly different than the idealized prediction
Central atom without lone pairs has the same real bond angle as the idealized angle.
The exceptions to this are square planar shapes and linear (derived from trigonal bipyramidal electronic structure) shapes where the lone pairs offset one another, thus causing no deviation from ideality.

19. Example
Lewis structure
Shape
Idealized bond angle
Real bond angle

20. Multiple Central Atoms
many molecules have larger structures with many interior atoms
we can think of them as having multiple central atoms
when this occurs, we describe the shape around each central atom in sequence

21. Representing 3-Dimensional Shapes on a 2-Dimensional Surface
one of the problems with drawing molecules is trying to show their dimensionality
by convention, the central atom is put in the plane of the paper
put as many other atoms as possible in the same plane and indicate with a straight line
for atoms in front of the plane, use a solid wedge
for atoms behind the plane, use a dashed wedge

22. Polarity of Molecules in order for a molecule to be polar it must
have polar bonds
electronegativity difference - theory
bond dipole moments - measured
have an unsymmetrical shape
vector addition
polarity affects the intermolecular forces of attraction
therefore boiling points and solubilities
like dissolves like
nonbonding pairs affect molecular polarity, strong pull in its direction

23. Molecule Polarity
The H-Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule.

25. Molecule Polarity
The O-C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule.

26. Molecule Polarity
The H-O bond is polar. The both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule.

27. Factors Affecting Dipole Moments
Lone-pair electrons on oxygen and nitrogen project out into space away from positively charged nuclei giving rise to a considerable charge separation and contributing to the dipole moment

28. Molecular Polarity Affects Solubility in Water
polar molecules are attracted to other polar molecules
since water is a polar molecule, other polar molecules dissolve well in water
and ionic compounds as well
some molecules have both polar and nonpolar parts

29. A Soap Molecule Sodium Stearate

30. Example - Decide Whether the Following Are PolarEN
O = 3.5
N = 3.0
Cl = 3.0
S = 2.5

31. Problems with Lewis Theory
Lewis theory gives good first approximations of the bond angles in molecules, but usually cannot be used to get the actual angle
Lewis theory cannot write one correct structure for many molecules where resonance is important
Lewis theory often does not predict the correct magnetic behavior of molecules
e.g., O2 is paramagnetic, though the Lewis structure predicts it is diamagnetic