1. Chapter 4.4: Intermolecular forces
The physical properties of molecular substances result from different types of forces between their molecules.

2. Important terms in this section
Intermolecular forces
Temporary/instantaneous dipole
Induced dipole
London (dispersion) forces
Permanent dipole
Dipole-dipole interaction
Van der Waal’s forces
Hydrogen bonding

3. Intermolecular vs Intramolecular forces
INTERmolecular forces: forces between molecules
Forces where whole molecules interact with each other
INTRAmolecular forces: forces within molecules
Forces that hold a single molecule together
So what are examples of intramolecular forces?
Examples of intermolecular forces?

4. Inter- and Intra-molecular forces
Just as the strength of the intramolecular forces can determine properties such as melting and boiling points, so can intermolecular forces
The increase in intermolecular forces will
Increase m.p. and b.p. (less volatile)
Can determine the Solubility (like dissolves like)
Conductivity (moving charges = conductivity)

5. London (dispersion) forces
Non-polar molecules can have:
Weak dipole = temporary/instantaneous dipole
This can influence a neighbor molecule = induced dipole

6. London (dispersion) forces
LDF strength increases as molecular size increases
More e- in atom = more chances for temporary dipoles occurring
LDF are ONLY forces between non-polar molecules
But exist in all molecules!!!
Generally have low m.p. and b.p. = Easy to break weak LDFs

7. Dipole-dipole attraction
Polar covalent molecules have a permanent dipole
These permanent charges orient other molecules the same way = dipole-dipole attraction
Strength of the attraction depends on distance & orientation of dipoles similar molecular mass

8. Van der Waal forces
Note: A polar molecule can induce a dipole in a non- polar molecule = dipole-induced dipole
Van der Waal forces: a term that includes both London dispersion forces and the dipole interactions (both dipole-dipole and dipole-induced dipole)
Note: this term does not include H-bonding!

9. Hydrogen bonding (H-bonding)
Is a special case of dipole-dipole attraction
Hydrogen bonded to:
Nitrogen
Fluorine
Oxygen
These are the strongest form of intermolecular attraction (so b.p. and m.p. are much higher than predicted)

10. H-bonding
Can you explain why water is less dense in solid than in liquid form?
H-bonds are the strongest intermolecular force, but is still weaker than covalent and ionic bonds

11. Intermolecular forces

12. Physical properties – M.P. & B.P.
What happens when CH3CH2OH is boiled?
H2O and CO2 are released as gases
CH3CH2OH is released as gas
CH3CH3 and O2 are released as gases
H2 and CO2 are released as gases
So are intermolecular forces broken or intramolecular forces broken when something is boiled?
Covalent cmpds have lower values than ionic
Weak intermolecular forces vs electrostatic attraction
Higher intermolecular bond strength, polar strength, and increase in molecular size increases m.p. and b.p.

13. Physical Properties – Solubility
“Like dissolves like”
Non-polar dissolve in non-polar (such as oil)
Due to the formation of LDF
Polar covalent molecules are usually soluble in water and other polar solvents
Dipole interactions and H-bonding
The larger the molecule with non-polar areas, the less polar the overall molecule
Giant covalent are insoluble in nearly all solvents
Too much E needed to break the bonds for interaction

14. Physical Properties – Electrical conductivity
Covalent compounds do not conduct as there are no charges particles to carry a charge
However, some polar covalent will conduct when in aqueous solution
Giant covalent molecules
Diamond = no conduction
Fullerene & silicon = semiconductors
Graphite & graphene = conductors

16. Chapter 4.5: Metallic Bonding
Metallic bonds involve a lattice of cations with delocalized electrons.

17. Delocalized electrons in metals
Metals have:
Few valence electrons (typically)
Low ionization energy
Form positive ions by losing electrons when reacted
When in elemental state, who can accept the e- s?
These electrons tend to “wander off” or be delocalized
These e- are no longer associated with 1 atom!

18. Metallic bond strength
The strength is determined by:
Number of delocalized e- (more e-, stronger bond)
Charge on cation (higher charge, stronger bond)
Radius of the cation (smaller radius, stronger bond)
Write the full e- configuration for Na and Mg:
Compare the above points and determine which will have stronger bonding. Why?

19. Metallic bonding strengths
Melting point of Na = 98 °C
Melting point of Mg = 650 °C
Which atom will have a higher melting point:
Na, K or Rb?
How did you determine this?
Melting point of K = 63 °C
Melting point of Rb = 39 °C
Transition metals: very strong metallic bonds. Why?
Many delocalized e- from 3d and 4s sub-shells

21. Alloys Alloys: Typically solid but made when molten
Solutions of metals
Enhanced properties
Typically solid but made when molten
Metallic bonds can accommodate other cations of different sizes into lattice
Bound by the delocalized e- in metal (thus metallically bonded)
Have different properties than original components
More chemically stable
Often stronger
Often more resistant to corrosion

22. Examples of Alloys