INTERMOLECULAR BONDING – VAN DER WAALS FORCES
(a) Van Der Waals Forces
Intermolecular attractions are attractions between one molecule and a neighbouring molecule.
The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as intramolecular attractions.
All molecules experience intermolecular attractions, although in some cases those attractions are very weak.
Even in a gas like hydrogen, H2, if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid.
In hydrogen’s case the attractions are so weak that the molecules have to be cooled to (-252°C) before the attractions are enough to condense the hydrogen as a liquid.
Helium’s intermolecular attractions are even weaker – the molecules won’t stick together to form a liquid until the temperature drops to (-269°C).
Polar molecules, such as water molecules, have a weak, partial negative charge at one region of the molecule (the oxygen atom in water) and a partial positive charge elsewhere (the hydrogen atoms in water).
Thus when water molecules are close together, their positive and negative regions are attracted to the oppositely-charged regions of nearby molecules.
The force of attraction, shown here as a dotted line, is called a hydrogen bond. Each water molecule is hydrogen bonded to four others.
The hydrogen bonds that form between water molecules account for some of the essential — and unique — properties of water.
- The attraction created by hydrogen bonds keeps water liquid over a wider range of temperature than is found for any other molecule its size.
- The energy required to break multiple hydrogen bonds causes water to have a high heat of vaporization; that is; a large amount of energy is needed to convert liquid water, where the molecules are attracted through their hydrogen bonds, to water vapor, where they are not.
Liquid Water and Hydrogen Bonding
Why water is a liquid?
In many ways, water is a miracle liquid. Since the hydrogen and oxygen atoms in the molecule carry opposite (though partial) charges, nearby water molecules are attracted to each other like tiny little magnets.
Hydrogen bonding makes water molecules “stick” together.
This makes water have high melting and boiling points compared to other covalent compounds such as ammonia (NH3) which have similar molecular mass but are gases
Ice and Hydrogen Bonding
The structure that forms in the solid ice crystal actually has large holes in it.
Therefore, in a given volume of ice, there are fewer water molecules than in the same volume of liquid water.
In other words, ice is less dense than liquid water and will float on the surface of the liquid.
Surface Tension and hydrogen bonding
As we just discussed, neighboring water molecules are attracted to one another.
Molecules at the surface of liquid water have fewer neighbors and, as a result, have a greater attraction to the few water molecules that are nearby.
This enhanced attraction is called surface tension.
It makes the surface of the liquid slightly more difficult to break through than the interior.
Water as a Solvent
The partial charge that develops across the water molecule helps make it an excellent solvent.
Water dissolves many substances by surrounding charged particles and “pulling” them into solution.
For example, common table salt, sodium chloride, is an ionic substance that contains alternating sodium and chlorine ions.
When table salt is added to water, the partial charges on the water molecule are attracted to the Na+ and Cl- ions.
Why does ethanol have a higher boiling point than methoxymethane?
Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-CH3, both have the same molecular formula, C2H6O.
They have the same number of electrons, and a similar length to the molecule.
The van der Waals attractions (both dispersion forces and dipole-dipole attractions) in each will be much the same.
However, ethanol has a hydrogen atom attached directly to oxygen – and that oxygen still has exactly the same two lone pairs as in a water molecule.
Hydrogen bonding can occur between ethanol molecules, although not as effectively as in water.
The hydrogen bonding is limited by the fact that there is only one hydrogen in each ethanol molecule with sufficient d+ charge.
In methoxymethane, the lone pairs on the oxygen are still there, but the hydrogens aren’t sufficiently d+ for hydrogen bonds to form.
Except in some rather unusual cases, the hydrogen atom has to be attached directly to the very electronegative element for hydrogen bonding to occur.
The boiling points of ethanol and methoxymethane show the dramatic effect that the hydrogen bonding has on the stickiness of the ethanol molecules: ethanol (with hydrogen bonding) 78.5°C methoxymethane (without hydrogen bonding) -24.8°C The hydrogen bonding in the ethanol has lifted its boiling point about 100°C.
It is important to realize that hydrogen bonding exists in addition to van der Waals attractions.
For example, all the following molecules contain the same number of electrons, and the first two are much the same length.
The higher boiling point of the butan-1-ol is due to the additional hydrogen bonding.