1. Polarity and Intermolecular (IM) Forces (Interactions)
Dr. Hima Joshi
Unless otherwise noted, images from Olmsted and Williams

2. Why do chemists think about intermolecular (IM) interactions?
Interactions between molecules can lead to chemical reactions.
Understanding interactions helps us predict solubility.
Take advantage of differences in solubility in the lab for separation of compounds.
Biochemical application: Since the bloodstream is mostly water, most drugs (made in labs) need to be water-soluble.

3. To understand intermolecular interactions, we need to better understand (intramolecular) bonding.
Bonding Continuum:
electrostatic
attraction
(NO sharing
of electrons)
equal sharing
of bonding pair
of electrons
unequal sharing
of bonding pair
of electrons
Figure: 01_05-10UN.jpg
Title:
Covalent, Polar Covalent, and Ionic Bonds
Caption:
While in a covalent bond the electrons are shared equally between the atoms, in a polar covalent bond, the shared electrons are held close to the most electronegative atom.
Notes:
In polar covalent compounds there are partial negative and partial positive charges on the atoms involved in the bond.
Analogy: tug of war
Cl has a better ability to pull bonding electrons toward itself
(as compared with C).

4. How do you know which atoms pull more on bonding electrons?
Electronegativity Scale
Electronegativity = ability to attract bonding
electrons
Helps you predict which bonds are polar.
(Helps predict unequal sharing.)

5. Electronegativity Scale
Electronegativity = ability to attract bonding electrons
• You will not be using the actual electronegativity values.
• You will use trends (from periodic table).
• F is the most electronegative.
• The closer to F on the periodic table, the more electronegative an element is.
• H belongs in same box as C. (No diff in electronegativity between H and C)
(So C-H is nonpolar.)
• To determine difference in electronegativity between two atoms, count how many steps each one is from F.
Figure: 01_07.jpg
Title:
Pauling Electronegativity Scale for Selected Elements
Caption:
The Pauling electronegativities of some of the elements found in organic compounds.
Notes:
Electronegativity is a guide for predicting the polar nature of bonds and the direction of their dipole moment. The more electronegative the atom, the more they “pull” the bonding electrons toward them. The most electronegative element on the periodic table is fluorine, with a Pauling electronegativity value of Although carbon has a slightly higher electronegativity than hydrogen, a C-H bond is considered to be nonpolar for all practical purposes.

6. Identify the polar covalent and nonpolar covalent bond(s)
in this molecule.
Figure: 01_06.jpg
Title:
Bond Polarity
Caption:
Bond polarity. Chloromethane contains a polar carbon-chlorine bond with a partial negative charge on chlorine and a partial positive charge on carbon. The electrostatic potential map shows a red region (electron-rich) around the partial negative charge and a blue region (electron-poor) around the partial positive charge. Other colors show intermediate values of electrostatic potential.
Notes:
electrostatic potential map
(red indicates slightly negative; blue is slightly positive)

7. In order for a molecule to be polar overall,
it must have individual bond dipoles that do not cancel out.
(The individual bond dipoles cannot be symmetrically
oriented in the molecule.)
Figure: 02_20-25UN.jpg
Title:
Electrostatic Potential Maps of Formaldehyde and CO2
Caption:
The electrostatic potential maps of formaldehyde and CO2 shows the bond dipole moments, with red at the negative ends and blue at the positive ends of the dipoles. In carbon dioxide the dipole moments are oriented in opposite directions, and they cancel each other.
Notes:
bond polarity vs. overall molecular polarity
(polar bond vs. polar molecule)

8. Determine if each of the following molecules is polar.
CH3Cl CHCl3 CCl4
Figure: 02_21.jpg
Title:
Molecular Dipole Moments
Caption:
Molecular dipole moments. A molecular dipole moment is the vector sum of the individual bond dipole moments.
Notes:

9. NH3 H2O (CH3)CO CH3CN Figure: 02_22.jpg Title: Effects of Lone Pairs
Caption:
The presence of lone pairs may have large effects on the molecular dipole moment.
Notes:

10. How would several polar molecules interact?
Dipole-dipole Interaction
Figure: 02_23.jpg
Title:
Dipole-dipole Interaction
Caption:
Dipole-dipole interactions result from the approach of two polar molecules. If their positive and negative ends approach, the interaction is an attractive one. If two negative ends or two positive ends approach, the interaction is repulsive. In a liquid or a solid, the molecules are mostly oriented with the positive and negative ends together, and the net force is attractive.
Notes:
Dipole-dipole interaction affects the physical properties of a compound. The greater the dipole-dipole interaction, the more energy will be required to vaporize or melt the compound, resulting in higher boiling and melting points.

11. Relative Strengths of Interactions
Type of Bond or
Interaction
Energy Required to
Break/Disrupt (kJ/mol)
Description
H-bond
10-40
Between slightly positive H and a lone pair
Ion-dipole
10-50
Between ion and
polar solvent
Dipole-dipole
3-4
Between polar
molecules
London Dispersion
1-10
Between all molecules
Note: 345 kJ/mol are required to break a C-C covalent bond.

12. What about molecules that don’t have dipole moments (are not polar)
What about molecules that don’t have dipole moments (are not polar)? What is their predominant intermolecular force/interaction?
Another example: two CO2 molecules

13. Relative Strengths of Interactions
Type of Bond or
Interaction
Energy Required to
Break/Disrupt (kJ/mol)
Description
H-bond
10-40
Between slightly positive H and a lone pair
Ion-dipole
10-50
Between ion and
polar solvent
Dipole-dipole
3-4
Between polar
molecules
London Dispersion
1-10
Between all molecules
Note: 345 kJ/mol are required to break a C-C covalent bond.

14. The electrons in a molecule are like a big, fluffy cloud.
This cloud can be temporarily
distorted, leading to an uneven
distribution of electron density
(partial charges).
London dispersion
(induced dipole)
forces
Figure: 02_24.jpg
Title:
London Dispersion Forces
Caption:
London dispersion forces result from the attraction of correlated temporary dipole moments.
Notes:
A temporary dipole moment in a molecule can induce a temporary dipole moment in a nearby molecule. An attractive dipole-dipole interactive results for a fraction of a second.

15. Hydrogen Bonding Hydrogen bond is not a covalent bond.
Figure: 02_25.jpg
Title:
Hydrogen Bonding
Caption:
Hydrogen bonding is a strong intermolecular attraction between an electrophilic O-H or N-H hydrogen atom and a pair of nonbonding electrons.
Notes:
Hydrogen bonding is an intermolecular interaction found in compounds with N-H and O-H bonds. The strong polar bond between the hydrogen and the heteroatom causes the hydrogen to interact with the lone pair(s) of the heteroatom on nearby molecules. The presence of hydrogen bonding will increase the boiling point of the compound because more energy will be required to break this interaction and vaporize the compound. Hydrogen bonding of O-H is stronger than the hydrogen bonding of N-H.
Hydrogen bond is not a covalent bond.
It is an interaction between a slightly positive H and a lone pair.
What makes an H slightly positive?
Identify some hydrogens in the structures above that CANNOT hydrogen bond.

16. Relative Strengths of Interactions
Type of Bond or
Interaction
Energy Required to
Break/Disrupt (kJ/mol)
Description
H-bond
10-40
Between slightly positive H and a lone pair
Ion-dipole
10-50
Between ion and
polar solvent
Dipole-dipole
3-4
Between polar
molecules
London Dispersion
1-10
Between all molecules
Note: 345 kJ/mol are required to break a C-C covalent bond.

17. H-bonding interactions are responsible for maintaining the structures of many biomolecules.
hemoglobin
DNA

18. Solubility: Like Dissolves Like
• Chemical species that can interact via dipole-dipole, H-bonding or ion-dipole interactions tend to mix well with (dissolve in) each other.
• Species with enough charges or partial charges tend to mix with (dissolve in) each other.
• Species without a significant number of charges (partial or actual) mix/dissolve in each other.

19. Why Oil and Water Don’t Mix
Are most of the bonds in triolein
polar or nonpolar?
What about the bonds in water?
triolein:
main component
of olive oil

20. Ion-dipole interaction Figure: 02_26.jpg Title:
Polar Solute in a Polar Solvent Dissolves
Caption:
The hydration of sodium and chloride ions by water molecules overcomes the lattice energy of sodium chloride. The salt dissolves.
Notes:
The water molecules will surround the sodium and chloride ions effectively dissolving them.
Ion-dipole
interaction

21. Relative Strengths of Interactions
Type of Bond or
Interaction
Energy Required to
Break/Disrupt (kJ/mol)
Description
H-bond
10-40
Between slightly positive H and a lone pair
Ion-dipole
10-50
Between ion and
polar solvent
Dipole-dipole
3-4
Between polar
molecules
London Dispersion
1-10
Between all molecules
Note: 345 kJ/mol are required to break a C-C covalent bond.

22. How do ion-dipole interactions apply to medicine?
Zoloft as a neutral molecule is not
soluble in water (bloodstream).
So, the hydrochloride salt of
Zoloft is used.
Turning a molecule into an ion
makes it soluble in water.
Zoloft

23. Acid-Base Chemistry and Solubility
+ NaHCO3

24. How do IM Interactions apply to phase changes?
Chemical Formula
Molecular Mass (amu)
Boiling Point
(oC)
H2O
~ 18
100
CH4
~ 16
– 161
When covalent molecules melt or boil, the bonds within
them do not break.
It is the interactions between them (intermolecular interactions)
that are disrupted.