INTERMOLECULAR FORCES in COVALENT COMPOUNDS
There are basically three physical states for a covalent compound, namely solid, liquid and gas. For an ionic compound it generally exist as a crystal.
The difference between the three states is the densities of the compound. Solid has the highest density, then the liquid, and finally the lowest density gaseous state. Since the density is a measure of how closely the molecules are packed per unit volume, we can conclude that it is a reflection of the attractive forces holding the molecules together.
Tutorial 1
What is the difference between a physical state and a phase? Answer
There are two classes of intermolecular forces between covalent molecules. The first type originated from polar molecules. These are easier to understand and we term them the Van der Waals forces. However even non-polar molecules must have attractive forces between them, otherwise they will all be gases which is not the case. The attraction between non-polar molecules is explained by London forces (after Fritz London).
LONDON INTERACTION
The first principle to remember is molecules are not in hibernation; there are many activities within them. For one the electrons will never stop moving. It will move around the atom and then across bonds to the other atoms. It moves to-and-flo taking the path of least resistance (in terms of energy barrier).
Because of such electron motion, there will be instantaneous transient charge imbalance. You can imagine that the dipole switch on and off so fast that on the average the molecule has no dipole. Such "split second dipoles" has a domino effect, as the magnetic field will cause its immediate neighbour to be polarised. The result is a dipole-induced dipole interaction.
The strength of the London interaction depends on how strongly the electrons are under the control of the nucleus. The larger the atom the further the electron orbitals are away from the nucleus and the electrons have a greater ability to roam freely. Such atoms have greater degree of instantaneous transient charge imbalance, and also a greater ease to have induced dipole.
Of course of all the intermolecular interactions the London interaction is the weakest. Examples will be all the interaction between the diatomic molecules - H2, N2, O2, F2, etc.
PHASE TRANSITION
Before we go on to other intermolecular attractions let us see how these forces are related to the physical state of matter.
The first principle is to remember that molecules are not in hibernation, they have their own intrinsic (personal) energy. One consequence of such energy is the molecule would like to be in a chaos state. They like to party. So if you like a cigarette you can see the fume floating through the room even though there is no air movement. This is often known as the translation (or Brownian) motion.
How freely the molecule can move will depend on the size of the particle and intermolecular attraction that exist. Small particles need less energy for translation motion, especially if the intermolecular attraction is only the London force.
At a temperature where the intermolecular interaction is greater than the energetic of the molecule the compound is a solid. As the temperature of the surrounding is increase, the energy of the molecule also increases. At a particular temperature the molecules break free from being "held tightly" to each other. It melts to a liquid. If you continue putting energy into the system the energy of the molecule increase further and it finally boils and become a gas.
| Particles | Melting Points / K | Boiling Points / K |
| Helium (He) | 1.8 | 4.2 |
| Hydrogen (H2) | 14.0 | 20.4 |
| Nitrogen (N2) | 63.2 | 77.4 |
| Oxygen (O2) | 54.4 | 90.2 |
| Argon (Ar) | 83.8 | 87.3 |
| Chlorine (Cl2) | 172.1 | 239.0 |
| Bromine (Br2) | 265.9 | 332.4 |
| Note: He and Ar were included to emphasise that the principles discussed is also true for the noble gas atoms. |
Tutorial 2
Answer
- What do you understand by melting and boiling?
- Guess the melting and boiling temperatures for neon, krypton, and methane.
- Explain the variation of melting point (M.Pt.) and boiling point (B.Pt.) of the following hydrocarbons.
Molecules Molecular Weight M.Pt./K B.Pt./K Methane CH4 16.0 90.7 111.7 Ethane CH3-CH3 30.1 89.9 184.6 Propane CH3-CH2-CH3 44.1 85.5 231.1 Butane CH3-[CH2]2-CH3 58.1 134.9 272.7 Pentane CH3-[CH2]3-CH3 72.1 143.5 309.3 Hexane CH3-[CH2]4-CH3 86.2 178.2 342.2
VAN DER WAALS INTERACTION
When different atoms bonded covalently to form molecules the sharing of electrons in the bond is never equal. It is influenced by the electronegativity of the atoms. Because of such unequal sharing of the two electrons in the bond there is always a slightly positive pole and a slightly negative pole along the bond. There exist a dipole for the bond. Such dipoles are not transient. It can be measured.
Since molecules can have many bonds the polar effect is the sum total of all the dipolar bonds. The geometry of the molecule of course has the greatest influence on the final polar property of the molecule.
Methane is a molecule with a carbon at the centre of a tetrahedral structure and the sigma C-H bond radiating from it at 109� from each other. The sum effect is that of a sphere with the hydrogen atoms on the surface of the sphere and the carbon atom embedded in the centre of the sphere. So there is no dipole for methane, although the surface will be slightly positively charged.
Tutorial 3
What polar properties would you expect from carbon tetrachloride? Answer
For chloroform CHCl3 molecule will possess a dipole. The surface of the sphere will have chloride atoms (more electronegative than carbon) at three points and a hydrogen atom (less electronegative than carbon) at one point.
Such dipoles will cause the molecules to be held onto each other more strongly than that with London forces. Even for uni-polar molecules like methane the induced-dipole-to-induced-dipole attraction is much stronger than that from the London forces.
So comparatively molecules held by Van der Waals attraction will have higher melting and boiling point as compared with molecule held together by London forces; for molecules of about the same size. The last part of the statement is important because the ability of the molecule to move around is also dependent on the size.
| Molecules | Molecular Weight | Melting Points | Boiling Points |
| Methane | 16.0 | -182.5�C | -161.4�C |
| Methyl chloride CH3Cl | 50.5 | -97.1�C | 24.2�C |
| Methylene chloride CH2Cl2 | 84.9 | -96.7�C | 39.8�C |
| Chloroform CHCl3 | 119.4 | -63.7�C | 61.7�C |
| Carbon tetrachloride CCl4 | 153.8 | -22.9�C | 76.7�C |
| Molecules | Molecular Weight | Melting Points | Boiling Points |
| CH3-CH3 | 30.1 | -183.3�C | -88.6�C |
| CH3-CH2Cl | 64.5 | -136.4�C | 12.3�C |
| CH3-CHCl2 | 99.0 | -97.4�C | 57.3�C |
| CH3-CCl3 | 133.4 | -32.6�C | 74.1�C |
HYDROGEN BONDING
For molecules with:
then it will be held to each other by a hydrogen bond.
Because of the extreme inequality of electron sharing in the A-H the hydrogen atom experiences high electron deficiency. Since B has a lone electron pair it help the hydrogen atom out by extending it a temporary loan of its electron. You can visually it as two scuba divers sharing one tank of oxygen gas.
The attraction between H and B is more than electrostatic, it is a very weak covalent bond with strength of about 21 kJ mol‾� (a bond strength of a σ-bond is about 410 kJ mol‾�). This makes it the strongest kind of (physical) intermolecular attraction force. The London forces have an intensity of about 2 kJ mol‾�, and the van der Waals forces about 2.5 kJ mol‾�. Of course this are only general values. Example
Tutorial 4
I like the way you explain things. I just read lesson it was very interesting. I have a question. Why is carbon dioxide a gas at room temperature while water is a liquid at room temperature? Amanda Petersen, 12th grader, USA. Answer
