All substances have one of two structures: either they have a giant structure, where the particles all bonded to one another in a single continuous 3D arrangement, or they are individual separate molecules. Only the latter have forces between these individual molecules, which are referred to as intermolecular forces.
Compared with the bonds that run throughout a giant structure, these intermolecular forces are weak. In order to melt or boil a substance made of simple molecules we have to overcome these forces, and since only a small amount of energy is required, their melting and boiling points are generally low. The strength of the intermolecular forces is reflected in how high or low the melting and boiling points are.
There are three types of intermolecular forces that hold molecules together. In order of increasing strength they are:
1) Induced dipole-dipole interactions, also called London forces
2) Permanent dipole-dipole interactions
3) Hydrogen bonding
In some questions and texts you may see the term Van der Waals forces. This is likely to be a reference to induced dipole-dipole interactions, but the term is ambiguous since permanent dipole-dipole interactions are also Van der Waals forces.
To understand intermolecular forces, we are going to need to understand what dipoles are: A dipole is a separation of positive and negative charges. As a result of such charge separation, the negative end of one dipole will attract the positive end of another dipole and vice versa, in a similar way to more familiar magnetic dipoles.
Induced dipole-dipole interactions (London forces)
If there were no forces between molecules, it would take no energy to separate them. Everything would be a gas, and the melting and boiling points would be absolute zero (0 Kelvin, equal to -273°C). In reality all molecules have London Forces – even the monatomic Noble Gases.
We find that the more electrons in a molecule, the stronger the London Forces holding one molecule to another e.g. the boiling points of the Noble Gases increase going down the group.
Induced dipole-dipole interactions arise because electrons are always moving (quickly, randomly within the atomic orbitals). At any moment in time it is possible for more electrons to lie on one side of the nucleus than the other. When this happens an instantaneous dipole occurs, with the nucleus the positive end of the dipole and the unbalanced electrons the negative end.
This instantaneous dipole produces an induced dipole in neighboring atoms or molecules, and oppositely-charged ends of these two dipoles attract one another.
Molecules having the same number of electrons as each other, for example isomers, are referred to as isoelectronic. As a result we might expect them to have London forces of the same strength, but in reality the melting and boiling points of isoelectronic molecules vary. London forces are short-range forces: they get stronger the closer the molecules become. This means that differences how close molecules can get to one another and how they pack together will be reflected in their melting and boiling points.
We consider these factors in terms of (NOT the same as surface area!). The straighter and more linear a molecule is, the more of its surface area can be adjacent to another molecule, conversely the more spherical it is, the less of its surface area can be in contact with an adjacent molecule:
Permanent Dipole-Dipole Interactions
Polar molecules exhibit an additional form of intermolecular force, known as permanent dipole-dipole interactions. Evidence for polar molecules being attracted to an electrostatic charge, when non-polar molecules are not, comes from the deflection of flowing polar molecules by a charged plastic rod.
<photo of experiment>
A polar molecule is one which has a dipole ‘locked into’ the molecule because of the distribution of charge between the bonded atoms. To understand how polar molecules give rise to permanent dipole-dipole intermolecular forces, we need to understand what gives a molecule a permanent dipole:
Electronegativity and Polar Bonds
If both nuclei attract the bonding electrons equally, they occupy the space between the nuclei, and we have a perfect covalent bond e.g. Cl2 which is a non-polar bond.
If one bonded atom attracts the bonding electrons more strongly than the other, because it
is more electronegative, than we have a polar bond – the electrons spend more time closer to the atom they are more attracted to: e.g. H2O, HCl, NH3
The symbols δ+ and δ- mean a partial positive and a partial negative charge – not as much as H+ or Cl– would be. We say that the bond has a dipole, and is therefore a polar bond.
Definition: Electronegativity is the ability of a bonded atom to attract the electrons in a covalent bond.
The more electronegative an atom is, the more it attracts the bonding electrons. It is therefore the DIFFERENCE in electronegativity between the two atoms in the bond which determines how polar the bond will be. Ionic and covalent are just the two extremes, with polar covalent bonds in between.
small difference/no difference: pure covalent bond e.g. C-H
moderate difference: polar (covalent) bond e.g. C-Cl
very large difference: ionic bond e.g. NaCl
We need to know how electronegative each of the atoms is in the bond to work out how polarized the bond will be, and what the resulting properties of the molecule will be.
Trends in electronegativity:
- The “top three” most electronegative elements are F > O > N
- Electronegativity increases across a period from Group 1 to Group 17
- Electronegativity decreases going down a group as atomic radius increases
- The electronegativity of hydrogen is lower than most non metals and higher than most metals.
Consequences for types of bond:
Potassium has an electronegativity of 0.8 while fluorine has an electronegativity of 4.0 so there is a very large difference in electronegativity. In potassium fluoride the bonding electrons essentially reside on the fluorine atom, so the bonding is ionic with very little covalent character.
Hydrogen has an electronegativity of 2.1 while oxygen has an electronegativity of 3.5 so there is a moderate difference in electronegativity. In water the bonds are covalent but the bonding electrons are attracted more towards the oxygen atom, so the bonds are polar.
Carbon has an electronegativity of 2.5 while hydrogen has an electronegativity of 2.1 so there is very little difference. The bonds in methane are covalent and the electrons are evenly shared between C and H atoms, so the bonds are non-polar.
Molecules with permanent dipoles (polar molecules)
We can use electronegativity to work out which bonds are polarized in a molecule. Whether the molecule has a permanent dipole then depends on whether the dipoles in the polar bonds balance each other out, given the 3-D shape of the molecule:
CH4 no polar bonds; no permanent dipole, not a polar molecule
CF4 has polar bonds, but symmetrically arranged; no permanent dipole; not a polar molecule
CH2O polar bond between C=O, C-H bonds not polar; has a permanent dipole; is a polar molecule
So molecules with polar bonds are not necessarily polar molecules, but polar molecules arise from unsymmetrical polar bonds. Consider CO and H2O as simple examples. H2O has an overall dipole but CO2 does not because the two dipoles cancel each other out.
Attraction between polar molecules
Polar molecules tend to have their permanent dipoles aligned, because the negative end of one permanent dipole attracts the positive end of the dipole in an adjacent molecule. This attraction is the permanent dipole-dipole interaction: a second intermolecular force in addition to London forces. Polar molecules therefore have higher melting and boiling points than similar non-polar molecules.
N2 63 non-polar molecule
O2 55 non-polar molecule
NO 110 polar molecule
Hydrogen bonding is not bonding in the normal sense but an intermolecular force. It is the strongest of all the intermolecular forces and results in some unexpected properties in substances that have it. A lot of biochemical systems/living systems depend on hydrogen bonding.
Evidence for a significantly stronger intermolecular force than can be explained by permanent dipole-dipole interactions and London forces comes from comparing the boiling points of the hydrides of Group 6. The trend shown by H2Te, H2Se and H2S would suggest that H2O should have a much lower boiling point than it actually does: water has hydrogen bonding while the other hydrides do not.
When do we get hydrogen bonds ?
Hydrogen bonds are a special case of permanent dipole-dipole interactions. They occur between the lone pair on a very electronegative atom and a hydrogen atom which is δ+ because it is bonded to a very electronegative atom.
Rules for when hydrogen bonding occurs between two molecules:
- One molecule has a H-atom which is very highly positively polarized
- The other molecule has one of the very electronegative atoms – fluorine, oxygen or nitrogen; and this atom has a lone pair available.
BOTH OF THESE CRITERIA MUST BE MET TO GET A HYDROGEN BOND
e.g. NH3 hydrogen bonding can occur between molecules
SiH4 no hydrogen bonding (no F, O or N atom)
HF hydrogen bonding can occur between molecules
HCl no hydrogen bonding (no F, O or N atom)
CH2O no hydrogen bonding (H is bonded to C, so not δ+)
Drawing a hydrogen bond between two molecules:
- Draw the two molecules, show lone pairs
- Label the δ+ and δ- atoms
- Show a H-bond from lone pair to δ+ H atom
- Make hydrogen bond angle 180°
As they are much stronger than the other types of intermolecular forces, substances with hydrogen bonding require much more energy to break the hydrogen bonds, and therefore have much higher melting and boiling points than similar molecules without hydrogen bonding.
Comparing the extent of hydrogen bonding:
If we have two different compounds both of which are capable of forming hydrogen bonds, we can determine which will have the strongest hydrogen bonding. (This will be reflected in that compound having a higher melting and boiling point).
- The more electronegative the δ- atom is, the stronger the hydrogen bond it forms will be.
- The more hydrogen bonds a molecule can form (count the available lone pairs on the electronegative δ- atoms), the stronger the hydrogen bonding will be.
e.g. H2O has stronger hydrogen bonding than NH3, because i) oxygen is more electronegative than nitrogen so the hydrogen bonds are stronger and ii) oxygen has two lone pairs so can form two hydrogen bonds whereas nitrogen has one lone pair and can only form one hydrogen bond. The hydrogen bonds in water take more energy to overcome than in ammonia, so water has a higher melting and boiling point than ammonia.
Predicting relative melting/boiling points
We are often asked to compare the predicted boiling or melting points of molecules in order to put them into order e.g. lowest to highest. To do this we have to first identify which types of intermolecular forces each molecule has.
- Write a table showing each molecule and which intermolecular forces it has
- For any molecules having only London forces, put these in order of number of electrons in the molecule (least electrons = weakest London forces)
- For molecules having only London forces and the same number of electrons, compare shapes and put these in order of area of surface contact (least area of surface contact = weakest London forces)
- Next will come any molecules with Permanent dipole-dipole interactions as well as London forces. We don’t have a simple basis for assessing the strength of the permanent dipole-dipole interactions, but we could use the strength of the London forces to order these molecules too
- Finally include the molecules with hydrogen bonding (ability to form more hydrogen bonds and having more electronegative atom with lone pair = stronger hydrogen bonding)
e.g. Put these molecules in order of increasing boiling point: CH3Cl, CCl4, CH3OH, CH4.
Molecule London Forces Permanent D-D Hydrogen bond
CH3Cl Yes Yes (polar C-Cl) No
CCl4 Yes No (symmetrical C-Cl) No
CH3OH Yes Yes Yes
CH4 Yes No (no polar bonds) No
Lowest boiling points will be CH4 then CCl4 (since CCl4 has more electrons than CH4). Next will come CH3Cl because that also has permanent dipole-dipole interactions, then finally CH3OH since this has hydrogen bonding as well.
CH4, CCl4, CH3Cl, CH3OH