1. Intermolecular forces
12/10/99
Questions
Why do some solids dissolve in water but others do not?
Why are some substances gases at room temperature, but others are liquid or solid?
What gives metals the ability to conduct electricity, what makes non-metals brittle?
The answers have to do with …
Intermolecular forces

2. Intermolecular Forces
Distinction between intramolecular bonding and intermolecular forces.
Intermolecular forces: van der Waals’, dipole-dipole, hydrogen bonding.
Effect of the intermolecular forces on the boiling point of a covalent substance.
Comparison of the boiling points of H2 and O2, of C2H4 and HCHO, and of H2O and H2S

3. Intermolecular forces
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Intermolecular forces
Overview
There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces
We have already looked at intramolecular bonds (ionic, polar, non-polar)
Intermolecular forces (IMF) have to do with the attraction between molecules
IMFs come in 3 flavours: 1dipole - dipole, ) H-bonding, 3) van der waals

4. A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary.
2 Phases in Glass
Solid phase - ice
Liquid phase - water

5. Intermolecular Forces
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
e.g. ionic or covalent bonds
Intermolecular vs Intramolecular
41 kJ to vaporize 1 mole of water (inter)
930 kJ to break all O-H bonds in 1 mole of water (intra)
“Measure” of intermolecular force
= DH
Generally, intermolecular forces are much weaker than intramolecular forces.

6. Intermolecular Forces
1. Dipole-Dipole Forces
Attractive forces between polar molecules and ionic substances
Orientation of Polar Molecules in a Solid
11.2

7. Ionic, Dipole - Dipole attractions
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We have seen that molecules can have a separation of charge
This happens in both ionic and polar bonds (the greater the EN, the greater the dipoles) H Cl + –
Molecules are attracted to each other in a compound by these +ve and -ve forces + – + – + – + –

8. Intermolecular Forces
Dipole-Dipole Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
11.2

9. 11.2

10. Intermolecular Forces
3. Hydrogen Bond
The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B or
A & B are N, O, or F
11.2

11. H - bonding
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H-bonding is a special type of dipole - dipole attraction that is very strong
It occurs when N, O, or F are bonded to H
Q- Calculate the EN for HCl and H2O
A- HCl = = 0.8, H2O = = 1.4
The high EN diff of NH, OH, and HF bonds cause these to be strong forces (about 5x stronger than normal dipole-dipole forces)
They are given a special name (H-bonding) because compounds containing these bonds are important in biological systems

13. Why is the hydrogen bond considered a “special” dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point
11.2

14. Van der waals forces
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Non-polar molecules do not have dipoles like polar molecules.
How, then, can non-polar compounds form solids or liquids?
These forces are due to small dipoles that exist in non-polar molecules
Because electrons are moving around in atoms there will be instants when the charge around an atom is not symmetrical
The resulting tiny dipoles cause attractions between atoms/molecules

15. Intermolecular Forces
2. Van der Walls (London Dispersion Forces)
Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules
No dipole
ion-induced dipole interaction
dipole-induced dipole interaction
11.2

16. Van der waalsforces Instantaneous dipole: Induced dipole:
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Instantaneous dipole:
Induced dipole:
Eventually electrons are situated so that tiny dipoles form
A dipole forms in one atom or molecule, inducing a dipole in the other

17. Intermolecular Forces
Van der Walls Continued
Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted.
Polarizability increases with:
greater number of electrons
more diffuse electron cloud
So Van der Walls forces usually increase with molar mass.
11.2

18. Van der Walls Forces Are present in all compounds
Can occur between atoms or molecules
Are due to electron movement not to EN
Are transient in nature (dipole-dipole are more permanent).
are weaker than dipole-dipole and H-bonding

19. What type(s) of intermolecular forces exist between each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: Wan der Walls forces. S O
SO2
SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.
11.2

20. Testing concepts
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Which attractions are stronger: intermolecular or intramolecular?
Intramolecular are stronger. A covalent bond is 100x stronger.
2. What evidence is there that nonpolar molecules attract each other?
Nonpolar molecules gather together as liquids or solids at low temperatures.
3. Which would have a lower boiling point: O2 or F2? Explain.
Both have Van der Walls forces but F2 would be lower because it is smaller. Larger atoms/molecules can have their electron clouds more easily deformed and thus have stronger Van der Walls attractions and higher melting/boiling points.
H2S has the strongest (-61).
4. Which would have a lower boiling point: NO or O2? Explain.
O2 because it has only Van der Walls forces. NO has a small EN, giving it small dipoles.

21. 1) a large EN 2) the small sizes of atoms.
5. Which would you expect to have the higher melting point (or boiling point): C8H18 or C4H10? Explain.
C8H18 would have the higher melting/boiling point. This is a result of the many more sites available for Van der Walls forces to form.
6. What two factors causes hydrogen bonds to be so much stronger than typical dipole-dipole bonds?
1) a large EN 2) the small sizes of atoms.
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22. 9. What kind(s) of intermolecular forces are present in the following substances:
a) NH3, b) SF6, c) PCl3, d) LiCl, e) HBr, f) CO2
(hint: consider EN and molecular shape/polarity)
a) NH3: Hydrogen bonding (H + N), Van der Walls
b) SF6: Van der Walls only (it is symmetrical).
c) PCl3: EN= Dipole-dipole, Van der Walls.
d) LiCl: EN= Ionic, (Van der Walls).
e) HBr: EN= Dipole-dipole, Van der Walls.
f) CO2: Van der Walls only (it is symmetrical)

23. Challenge: Ethanol (CH3CH2OH) and dimethyl ether (CH3OCH3) have the same formula (C2H6O). Ethanol boils at 78 C, whereas dimethyl ether boils at -24 C. Explain why the boiling point of the ether is so much lower than the boiling point of ethanol.
Challenge: In ethanol, H and O are bonded (the large EN results in H-bonding). In dimethyl ether the O is bonded to C (a smaller EN results in a dipole-dipole attraction rather than hydrogen bonding).

24. H – bonding and boiling point
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See pg. 369 – Q – why does BP as period , why are some BP high at period 2?

25. For more lessons, visit www.chalkbored.com
Graph Q
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Boiling points increase down a group (as period increases) for two reasons: 1) EN tends to increase and 2) size increases. A larger size means greater London forces.
Boiling points are very high for H2O, HF, and NH3 because these are hydrogen bonds (high EN), creating large intermolecular forces
For more lessons, visit

26. 12/10/99
Testing concepts
A) F2 would be lower because it is smaller. Larger atoms/molecules can have their electron clouds more easily deformed and thus have stronger London attractions and higher melting/boiling points.
B) O2 because it has only London forces. NO has a small EN, giving it small dipoles.

27. Testing concepts a) NH3: Hydrogen bonding (H + N), London.
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Testing concepts
a) NH3: Hydrogen bonding (H + N), London.
b) SF6: London only (it is symmetrical).
c) PCl3: EN= Dipole-dipole, London.
d) LiCl: EN= Ionic, (London).
e) HBr: EN= Dipole-dipole, London.
f) CO2: London only (it is symmetrical)