Intermolecular Forces of Attraction

Intermolecular Forces of Attraction

METALLIC BONDING

What is a metallic bond?

Metallic bonding in sodium

Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms. Even a metal like sodium (melting point 97.8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the Periodic Table.

Sodium has the electronic structure 1s22s22p63s1. When sodium atoms come together, the electron in the 3s atomic orbital of one sodium atom shares space with the corresponding electron on a neighbouring atom to form a molecular orbital - in much the same sort of way that a covalent bond is formed.

The difference, however, is that each sodium atom is being touched by eight other sodium atoms - and the sharing occurs between the central atom and the 3s orbitals on all of the eight other atoms. And each of these eight is in turn being touched by eight sodium atoms, which in turn are touched by eight atoms - and so on and so on, until you have taken in all the atoms in that lump of sodium.

All of the 3s orbitals on all of the atoms overlap to give a vast number of molecular orbitals which extend over the whole piece of metal. There have to be huge numbers of molecular orbitals, of course, because any orbital can only hold two electrons.

The electrons can move freely within these molecular orbitals, and so each electron becomes detached from its parent atom. The electrons are said to be delocalized. The metal is held together by the strong forces of attraction between the positive nuclei and the delocalised electrons.

This is sometimes described as "an array of positive ions in a sea of electrons".

If you are going to use this view, beware! Is a metal made up of atoms or ions? It is made of atoms.

Each positive centre in the diagram represents all the rest of the atom apart from the outer electron, but that electron hasn't been lost - it may no longer have an attachment to a particular atom, but it's still there in the structure. Sodium metal is therefore written as Na - not Na+.

Metallic bonding in magnesium

If you work through the same argument with magnesium, you end up with stronger bonds and so a higher melting point.

Magnesium has the outer electronic structure 3s2. Both of these electrons become delocalized, so the "sea" has twice the electron density as it does in sodium. The remaining "ions" also have twice the charge (if you are going to use this particular view of the metal bond) and so there will be more attraction between "ions" and "sea".

More realistically, each magnesium atom has one more proton in the nucleus than a sodium atom has, and so not only will there be a greater number of delocalized electrons, but there will also be a greater attraction for them.

Magnesium atoms have a slightly smaller radius than sodium atoms, and so the delocalized electrons are closer to the nuclei. Each magnesium atom also has twelve near neighbours rather than sodium's eight. Both of these factors increase the strength of the bond still further.

Metallic bonding in transition elements

Transition metals tend to have particularly high melting points and boiling points. The reason is that they can involve the 3d electrons in the delocalization as well as the 4s. The more electrons you can involve, the stronger the attractions tend to be.

The metallic bond in molten metals

In a molten metal, the metallic bond is still present, although the ordered structure has been broken down. The metallic bond isn't fully broken until the metal boils. That means that boiling point is actually a better guide to the strength of the metallic bond than melting point is. On melting, the bond is loosened, not broken.

INTERMOLECULAR BONDING - VAN DER WAALS FORCES

What are intermolecular attractions?

Intermolecular versus intramolecular bonds

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. These two words are so confusingly similar that it is safer to abandon one of them and never use it. The term "intramolecular" won't be used again on this site.

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 21 K (-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 4 K (-269°C).

van der Waals forces: dispersion forces

Dispersion forces (one of the two types of van der Waals force we are dealing with on this page) are also known as "London forces" (named after Fritz London who first suggested how they might arise).

The origin of van der Waals dispersion forces

Temporary fluctuating dipoles

Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts. But that's only true on average.

The lozenge-shaped diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion.

But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end -. The other end will be temporarily short of electrons and so becomes +.

An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule.

This constant "sloshing around" of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom.

If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer properly covered by electrons for that instant.

How temporary dipoles give rise to intermolecular attractions

I'm going to use the same lozenge-shaped diagram now to represent any molecule which could, in fact, be a much more complicated shape. Shape does matter (see below), but keeping the shape simple makes it a lot easier to both draw the diagrams and understand what is going on.

Imagine a molecule which has a temporary polarity being approached by one which happens to be entirely non-polar just at that moment. (A pretty unlikely event, but it makes the diagrams much easier to draw! In reality, one of the molecules is likely to have a greater polarity than the other at that time - and so will be the dominant one.)

As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one.

This sets up an induced dipole in the approaching molecule, which is orientated in such a way that the + end of one is attracted to the - end of the other.

An instant later the electrons in the left hand molecule may well have moved up the other end. In doing so, they will repel the electrons in the right hand one.

The polarity of both molecules reverses, but you still have + attracting -. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronization so that the attraction is always maintained.

There is no reason why this has to be restricted to two molecules. As long as the molecules are close together this synchronized movement of the electrons can occur over huge numbers of molecules.

This diagram shows how a whole lattice of molecules could be held together in a solid using van der Waals dispersion forces. An instant later, of course, you would have to draw a quite different arrangement of the distribution of the electrons as they shifted around - but always in synchronization.

The strength of dispersion forces

Dispersion forces between molecules are much weaker than the covalent bonds within molecules. It isn't possible to give any exact value, because the size of the attraction varies considerably with the size of the molecule and its shape.

How molecular size affects the strength of the dispersion forces

The boiling points of the noble gases are

helium / / -269°C
neon / / -246°C
argon / / -186°C
krypton / / -152°C
xenon / / -108°C
radon / / -62°C

All of these elements have monatomic molecules.

The reason that the boiling points increase as you go down the group is that the number of electrons increases, and so also does the radius of the atom. The more electrons youhave, and the more distance over which they can move, the bigger the possible temporary dipoles and therefore the bigger the dispersion forces.

Because of the greater temporary dipoles, xenon molecules are "stickier" than neon molecules. Neon molecules will break away from each other at much lower temperatures than xenon molecules - hence neon has the lower boiling point.

This is the reason that (all other things being equal) bigger molecules have higher boiling points than small ones. Bigger molecules have more electrons and more distance over which temporary dipoles can develop - and so the bigger molecules are "stickier".

How molecular shape affects the strength of the dispersion forces

The shapes of the molecules also matter. Long thin molecules can develop bigger temporary dipoles due to electron movement than short fat ones containing the same numbers of electrons.

Long thin molecules can also lie closer together - these attractions are at their most effective if the molecules are really close.

For example, the hydrocarbon molecules butane and 2-methylpropane both have a molecular formula C4H10, but the atoms are arranged differently. In butane the carbon atoms are arranged in a single chain, but 2-methylpropane is a shorter chain with a branch.

Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules.

van der Waals forces: dipole-dipole interactions

A molecule like HCl has a permanent dipole because chlorine is more electronegative than hydrogen. These permanent, in-built dipoles will cause the molecules to attract each other rather more than they otherwise would if they had to rely only on dispersion forces.

It's important to realize that all molecules experience dispersion forces. Dipole-dipole interactions are not an alternative to dispersion forces - they occur in addition to them. Molecules which have permanent dipoles will therefore have boiling points rather higher than molecules which only have temporary fluctuating dipoles.

Surprisingly dipole-dipole attractions are fairly minor compared with dispersion forces, and their effect can only really be seen if you compare two molecules with the same number of electrons and the same size. For example, the boiling points of ethane, CH3CH3, and fluoromethane, CH3F, are:

Why choose these two molecules to compare? Both have identical numbers of electrons, and if you made models you would find that the sizes were similar - as you can see in the diagrams. That means that the dispersion forces in both molecules should be much the same.

The higher boiling point of fluoromethane is due to the large permanent dipole on the molecule because of the high electronegativity of fluorine. However, even given the large permanent polarity of the molecule, the boiling point has only been increased by some 10°.

Here is another example showing the dominance of the dispersion forces. Trichloromethane, CHCl3, is a highly polar molecule because of the electronegativity of the three chlorines. There will be quite strong dipole-dipole attractions between one molecule and its neighbours.

On the other hand, tetrachloromethane, CCl4, is non-polar. The outside of the molecule is uniformly - in all directions. CCl4 has to rely only on dispersion forces.

So which has the highest boiling point? CCl4 does, because it is a bigger molecule with more electrons. The increase in the dispersion forces more than compensates for the loss of dipole-dipole interactions.

The boiling points are:

CHCl3 / / 61.2°C
CCl4 / / 76.8°C

INTERMOLECULAR BONDING - HYDROGEN BONDS

The evidence for hydrogen bonding

Many elements form compounds with hydrogen - referred to as "hydrides". If you plot the boiling points of the hydrides of the Group 4 elements, you find that the boiling points increase as you go down the group.

The increase in boiling point happens because the molecules are getting larger with more electrons, and so van der Waals dispersion forces become greater.

If you repeat this exercise with the hydrides of elements in Groups 5, 6 and 7, something odd happens.

Although for the most part the trend is exactly the same as in group 4 (for exactly the same reasons), the boiling point of the hydride of the first element in each group is abnormally high.

In the cases of NH3, H2O and HF there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break. These relatively powerful intermolecular forces are described as hydrogen bonds.

The origin of hydrogen bonding

The molecules which have this extra bonding are:

Note: The solid line represents a bond in the plane of the screen or paper. Dotted bonds are going back into the screen or paper away from you, and wedge-shaped ones are coming out towards you.

Notice that in each of these molecules:

• The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge.
• Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair.

Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things.

Consider two water molecules coming close together.

The + hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a co-ordinate (dative covalent) bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction.

Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water. If you liken the covalent bond between the oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status. On the same scale, van der Waals attractions represent mere passing acquaintances!

Water as a "perfect" example of hydrogen bonding

Notice that each water molecule can potentially form four hydrogen bonds with surrounding water molecules. There are exactly the right numbers of + hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding.

This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride. In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen only has one lone pair. In a group of ammonia molecules, there aren't enough lone pairs to go around to satisfy all the hydrogens.

In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there is exactly the right number of each. Water could be considered as the "perfect" hydrogen bonded system.

More complex examples of hydrogen bonding

The hydration of negative ions

When an ionic substance dissolves in water, water molecules cluster around the separated ions. This process is called hydration.

Water frequently attaches to positive ions by co-ordinate (dative covalent) bonds. It bonds to negative ions using hydrogen bonds.

The diagram shows the potential hydrogen bonds formed to a chloride ion, Cl-. Although the lone pairs in the chloride ion are at the 3-level and wouldn't normally be active enough to form hydrogen bonds, in this case they are made more attractive by the full negative charge on the chlorine.

However complicated the negative ion, there will always be lone pairs that the hydrogen atoms from the water molecules can hydrogen bond to.

Hydrogen bonding in alcohols

An alcohol is an organic molecule containing an -O-H group.

Any molecule which has a hydrogen atom attached directly to an oxygen or a nitrogen is capable of hydrogen bonding. Such molecules will always have higher boiling points than similarly sized molecules which don't have an -O-H or an -N-H group. The hydrogen bonding makes the molecules "stickier", and more heat is necessary to separate them.