1. Intermolecular forces

2. Intermolecular forces
Remember this phrase . . .
Intermolecular forces

3. It’s the answer to (nearly) every question about physical properties.
Why is H2O a liquid, but H2S is a gas?
Intermolecular forces!
Why is ethanol a liquid but ethane is a gas?
Why wasn’t life on Earth wiped out during various ice ages?
Intermolecular forces are forces between molecules which are not chemically bonded.

4. We split intermolecular forces into 3 types.
The easiest to understand are
permanent dipole – permanent dipole forces
Sometimes we just call these dipole-dipole forces
These forces are about 1% as strong as a covalent bond.

5. If a molecule is a dipole (make sure you know how to tell!)
Then the +ve end of one molecule is attracted to the –ve end of another molecule.
This builds up a whole network of forces holding molecules in position.
Energy must be supplied to overcome these forces before a substance can melt or boil.
Polar molecules have higher boiling points than non-polar molecules

7. Hydrogen bonds

8. A special case of dipole-dipole force is hydrogen bonding.
A hydrogen bond is about 10% as strong as a covalent bond.

9. Hydrogen bonding happens when a very electronegative atom is bonded to a hydrogen atom.
ONLY N, O or F
The electronegative atom withdraws the bonding electrons
Hydrogen has no other (shielding) electrons
So the proton in the nucleus is exposed,
Meaning that any electrostatic forces are relatively large.

10. Hydrogen bonding in water

11. Hydrogen bonding in ammonia

12. Hydrogen bonding in HF

13. Hydrogen bonding in organic chemicals

14. Because hydrogen bonding is a stronger force than dipole-dipole attraction,
Molecules where hydrogen bonding is present have higher BPs

16. Why is ice less dense than water?
The molecules of H2O in water are bouncing randomly around, which makes it a liquid. Usually molecules move closer together when a liquid becomes a solid, but because H2O molecules are strongly polarized, when they slow down and become a solid, they arrange themselves into a matrix that actually pushes them further apart--making ice less dense than water.

17. Also . . .
What might happen to the bond angle as hydrogen bonds form?
The lone pair electrons are pulled slightly further away from the oxygen atoms
So they repel O-H bonds slightly less
And the bond angle increases slightly
So each water molecule takes up slightly more room.

19. Van der Waal’s Forces
The third kind of intermolecular force is known as:
Van der Waal’s forces
London Forces
Temporary dipole – induced dipole forces
These are all the same thing!

20. Van der Waal’s forces exist between all molecules – even ones which aren’t polar!
Molecules have lots of electrons which are in permanent random motion.
At certain times (just by chance) more of the electrons will be on one side of a molecule
And the molecule will have a temporary dipole

21. Temporary dipole

22. The temporary dipole attracts or repels the electrons in neighbouring molecules.
This creates an induced dipole

24. Van der Waal’s forces are very weak – less than 1% the strength of a covalent bond
They act over very short distances only
The electrons don’t stay still, so the dipoles change.

25. The strength of Van der Waal’s forces depends on the number of electrons present.
Bigger molecules (higher Mr) have more electrons, so they have greater VdW forces, and a higher BP.
The strength of Van der Waal’s forces also depends on the amount of surface area in contact with another molecule.
Long molecules have a higher BP than short or spherical molecules.

26. Explain why pentane has a BP of 309K, but dimethyl propane has a BP of 283K.
Short answer – intermolecular forces!
Long answer – not polar so only VdW’s forces.
Same Mr, so no difference in number of electrons.
Pentane is a straight chain molecule, so relatively large contact area.
Dimethyl propane has side chains, so takes a spherical shape.
Lower contact area, so weaker forces.