Three types of force can operate between covalent molecules:
- Dispersion Forces
also known as London Forces (named after Fritz London who first described these forces theoretically 1930) or as Weak Intermolecular Forces or as van der Waal's Forces (namd after the person who contributed to our understanding of non-ideal gas behaviour).
- Dipole-dipole interactions
- Hydrogen bonds
Relative strength of Intermolecular Forces:
- Intermolecular forces (dispersion forces, dipole-dipole interactions and hydrogen bonds) are much weaker than intramolecular forces (covalent bonds, ionic bonds or metallic bonds)
- dispersion forces are the weakest intermolecular force (one hundredth-one thousandth the strength of a covalent bond), hydrogen bonds are the strongest intermolecular force (about one-tenth the strength of a covalent bond).
- dispersion forces < dipole-dipole interactions < hydrogen bonds
Dispersion Forces (London Forces, Weak Intermolecular Forces, van der Waal's Forces)
- are very weak forces of attraction between molecules resulting from:
- momentary dipoles occurring due to uneven electron distributions in neighbouring molecules as they approach one another
- the weak residual attraction of the nuclei in one molecule for the electrons in a neighbouring molecule.
- The more electrons that are present in the molecule, the stronger the dispersion forces will be.
- Dispersion forces are the only type of intermolecular force operating between non-polar molecules, for example, dispersion forces operate between hydrogen (H2) molecules, chlorine (Cl2) molecules, carbon dioxide (CO2) molecules, dinitrogen tetroxide (N2O4) molecules and methane (CH4) molecules.
- are stronger intermolecular forces than Dispersion forces
- occur between molecules that have permanent net dipoles (polar molecules), for example, dipole-dipole interactions occur between SCl2 molecules, PCl3 molecules and CH3Cl molecules.
If the permanent net dipole within the polar molecules results from a covalent bond between a hydrogen atom and either fluorine, oxygen or nitrogen, the resulting intermolecular force is referred to as a hydrogen bond (see below).
- The partial positive charge on one molecule is electrostatically attracted to the partial negative charge on a neighbouring molecule.
- occur between molecules that have a permanent net dipole resulting from hydrogen being covalently bonded to either fluorine, oxygen or nitrogen. For example, hydrogen bonds operate between water (H2O) molecules, ammonia (NH3) molecules, hydrogen fluoride (HF) molecules, hydrogen peroxide (H2O2) molecules, alkanols (alcohols) such as methanol (CH3OH) molecules, and between alkanoic (caboxylic) acids such as ethanoic (acetic) acid (CH3COOH) and between organic amines such as methanamine (methyl amine, CH3NH2).
- are a stronger intermolecular force than either Dispersion forces or dipole-dipole interactions since the hydrogen nucleus is extremely small and positively charged and fluorine, oxygen and nitrogen being very electronegative so that the electron on the hydrogen atom is strongly attracted to the fluorine, oxygen or nitrogen atom, leaving a highly localised positive charge on the hydrogen atom and highly negative localised charge on the fluorine, oxygen or nitrogen atom. This means the electrostatic attraction between these molecules will be greater than for the polar molecules that do not have hydrogen covalently bonded to either fluorine, oxygen or nitrogen.
Effect of Intermolecular forces on melting and boiling points of molecular covalent substances:
Since melting or boiling result from a progressive weakening of the attractive forces between the covalent molecules, the stronger the intermolecular force is, the more energy is required to melt the solid or boil the liquid.
|If only dispersion forces are present, then the more electrons the molecule has (and consequently the more mass it has) the stronger the dispersion forces will be, so the higher the melting and boiling points will be.
Consider the hydrides of Group IV, all of which are non-polar molecules, so only dispersion forces act between the molecules.
CH4 (molecular mass ~ 16), SiH4 (molecular mass ~ 32), GeH4 (molecular mass ~ 77) and SnH4 (molecular mass ~ 123) can all be considered non-polar covalent molecules.
As the mass of the molecules increases, so does the strength of the dispersion force acting between the molecules, so more energy is required to weaken the attraction between the molecules resulting in higher boiling points.
|Boiling Points of Group IV Hydrides
|If a covalent molecule has a permanent net dipole then the force of attraction between these molecules will be stronger than if only dispersion forces were present between the molecules. As a consequence, this substance will have a higher melting or boiling point than similar molecules that are non-polar in nature.
Consider the boiling points of the hydrides of Group VII elements.
All of the molecules HF (molecular mass ~ 20), HCl (molecular mass ~ 37), HBr (molecular mass ~ 81) and HI (molecular mass ~ 128) are polar, the hydrogen atom having a partial positive charge (H) and the halogen atom having a partial negative charge (F, Cl, Br, I).
As a consequence, the stronger dipole-interactions acting between the hydride molecules of Group VII elements results in higher boiling points than for the hydrides of Group IV elements as seen above.
With the exception of HF, as the molecular mass increases, the boiling point of the hydrides increase.
HF is an exception because of the stronger force of attraction between HF molecules resulting from hydrogen bonds acting bewteen the HF molecules. Weaker dipole-dipole interactions act between the molecules of HCl, HBr and HI. So HF has a higher boiling point than the other molecules in this series.
|Boiling Points of Group VII hydrides
Effect of Intermolecular Forces on Solubility
In general like dissolves like: