1. Intermolecular Forces (4.3.1)

2. Dipole-Dipole Forces
Dipole-dipole forces exist between neutral polar molecules.
Polar molecules need to be close together.
There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble.
If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

3. Dipole-Dipole Forces

4. Dipole-Dipole Forces

5. London Dispersion (van der Waals) Forces
Weakest of all intermolecular forces.
It is possible for two adjacent neutral molecules to affect each other.
The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom).
For an instant, the electron clouds become distorted.
In that instant a dipole is formed (called an instantaneous/temporary dipole).

6. One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom).
The forces between instantaneous dipoles are called London dispersion forces.

7. Polarizability is the ease with which an electron cloud can be deformed.
The larger the molecule (the greater the number of electrons) the more polarizable.
London dispersion forces increase as molecular weight increases.
London dispersion forces exist between all molecules.
London dispersion forces depend on the shape of the molecule.

8. The greater the surface area available for contact, the greater the dispersion forces.
London dispersion forces between spherical molecules are lower than between sausage-like molecules.

9. London Dispersion Forces

10. London Dispersion Forces

11. Hydrogen Bonding
Special case of dipole-dipole forces.
By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high.
Intermolecular forces are abnormally strong.

12. H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N).
Electrons in the H-X (X = electronegative element) lie much closer to X than H.
H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.
Therefore, H-bonds are strong.

14. Hydrogen Bonding

15. Hydrogen bonds are responsible for:
Ice Floating
Solids are usually more closely packed than liquids;
Therefore, solids are more dense than liquids.
Ice is ordered with an open structure to optimize H-bonding.
Therefore, ice is less dense than water.
In water the H-O bond length is 1.0 Å.
The O…H hydrogen bond length is 1.8 Å.
Ice has water molecules arranged in an open, regular hexagon.
Each + H points towards a lone pair on O.

18. Summary of Intermolecular Forces
Ionic Compound – contains positive ions (usually metals) and negative ions. Held together by ionic bonds.
Covalent Compound – contains 2 or more nonmetals (no ions)
Nonpolar – contains only London dispersion forces (LDF)
Polar – contains LDF and dipole-dipole forces
Polar with H bonded to N, O, or F (with unshared pair) – contains LDF, dipole-dipole forces, and hydrogen bonds.
Larger molecule, stronger LDF (all other factors equal)
Mixture – ions mixed with polar molecules – contains ion-dipole forces (very strong) – like Na+ and Cl- ions in water.

19. Examples – determine the types of forces present in each:
H2O
CCl4
SO2
LiF
Ca(NO3)2 aqueous solution HF
PCl3

20. Examples – determine the types of forces present in each:
H2O LDF, dipole-dipole, H-bonds
CCl4 LDF
SO2 LDF and dipole-dipole
LiF ionic bonds
Ca(NO3)2 aqueous solution ion-dipole forces
HF LDF, dipole-dipole, H-bonds
PCl3 LDF and dipole-dipole

21. Relative Strengths of Forces
Bonds (covalent, ionic, metallic)
Hydrogen bonds
Dipole-dipole forces
London dispersion forces