Liquids, Solids and Changes of State
Physical Properties versus Chemical Properties
Although much chemistry is devoted to chemical properties and reactions, the physical properties of substances often concern us more on a day-to-day basis. Farmers worry about whether the expected precipitation will come in liquid or solid form. Rain, a liquid, is usually welcomed; hail, a solid, is almost universally dreaded because of the damage it can inflict on crops. The fact that water expands when it solidifies may bring disaster (or at least an expensive repair bill) to the automobile owner who has neglected to put antifreeze in the car's radiator if the temperature drops very far below freezing. During summer, the high rate of evaporation of paint solvent can make painting in the hot sun difficult. In Banff the traditional "three minute egg" isn't nearly as well cooked as most people like, so it's boiled a little longer.
These are just a few example that show how the physical properties of substances and the transformations among the three states of matter influence our lives. One of the goals of chemistry has been to understand what determines these properties.
Why Gases Differ From Liquids and Solids
Physical properties depend on forces of attraction between molecules, which are strong in liquids and solids and very weak in gases.
Most of the physical properties of gases, liquids, and solids are actually controlled by the strengths of intermolecular attractions - the attractive forces that exist between neighbouring particles. In liquids and solids, these forces are much stronger than in gases. But why?
Anyone who has ever played with magnets has discovered an important property of attractive forces between things. Whether these forces are magnetic or electrical, their strengths decrease quite rapidly as the distance between the attracting particles increases. For example, when two magnets are placed near each other, the attraction between them can be quite strong, but if they are far apart hardly any attraction is felt at all. This same phenomenon is what is primarily responsible for that fact that gases have properties that depend very little on chemical composition, whereas the properties of liquids and solids are influenced by chemical composition to a very large extent.
In a gas the molecules are far apart. At these large distances, the attractive forces are very weak, so the differences caused by dissimilarities in chemical makeup are very small and are hardly noticeable at all. As a result, all gases seem to be pretty much alike. However, in a liquid or solid, where the molecules are very close to each other, these attractive forces become quite strong. Differences between the intermolecular attractions caused by variations in chemical composition become magnified, and this causes unlike substances to behave quite differently.
Intermolecular Attractions
Dipole-dipole forces, London forces, and hydrogen bonding are the principle kinds of intermolecular attractions.
The attractions between molecules are always much weaker than the attractions within molecules. In an molecule of HCl, the hydrogen atom and chlorine atom are held to each other very tightly by a covalent bond. Neighbouring HCl molecules are attracted to each other by much weaker forces. Therefore, when a particular chlorine atom moves, the hydrogen atom bonded to it is forced to go along, and the HCl molecule remains intact as it moves about.
Dipole-Dipole Attractions
Dipole molecules like HCl tend to line up so that the positive end of one dipole is near the negative end of another. Thermal energy, however, causes the molecules to jiggle about, so the alignment isn't perfect. Nevertheless, there is still a net attraction between the polar molecules. These attractive forces are called dipole-dipole attractions. Generally they are only about 1% as strong as a covalent bond, and their strengths decrease quite rapidly as the distance between molecules increases.

London Forces
It is fairly easy to understand why there are attractions between polar molecules such as HCl or SO2. But even between the particles of nonpolar substances such as Cl2 or CH4, there are attractive forces. Both Cl2 and CH4 can be condensed to liquids and even solids if cooled to low enough temperatures, so there must be attractions between their particles that cause them to cling together in the liquid or solid state. These attractions between nonpolar particles are considerably weaker then those of polar molecules.
Take a look at the table below for two polar and two nonpolar substances.
Substance Boiling Point Formula weight Number of Electrons

HCl                -84.9                36                       18

H2S               -60.7                 34                       18


F2                    -188.1                 38                        18

Ar               -185.7                 40                         18

All four are somewhat similar in that they have similar formula weights and the same number of electrons. Notice that the two nonpolar substances have very similar boiling points. Notice also that the two polar compounds have similar boiling points. However, the nonpolar substances have boiling points that are at least 100 degrees lower then those of the polar substances, which tells us that the attractions in the nonpolar liquids are considerably weaker then those in the polar liquids. They only become noticeable when the thermal energy is low enough for the molecules to slow down so that the forces can work.
In 1930, Fritz London, a German physicist, offered a simple explanation for the weak attractions between nonpolar particles. In any atom or molecule the electrons are constantly moving, and if we could examine this motion in two neighbouring particles, we would find that the movement of electrons in one particle influences the movement of electrons in the other. 

This is because electrons repel each other and tend to stay as far apart as then can. Therefore, as an electron of one particle gets near the other particle, electrons on the second are pushed away. This happens continually as the electrons move around, so to some extent, the electron density in both particles flickers back and forth in a synchronous fashion. Take a look at the series of instantaneous "frozen" views below.

Notice that at any given moment the electron density of a particle can be unsymmetrical, with more negative charge on one side than on the other. For that particular instant, the particle becomes a dipole, and we call it a momentary dipole, or an instantaneous dipole.
As the instantaneous dipole forms in one particle, it pushes electron density around in its neighbour, which causes the electron density there to become unsymmetrical, too. As a result, this second particle also becomes a dipole, and we call it an induced dipole because it is caused by, or induced by, the formation of the first dipole. Since the electrons are in constant motion these dipoles flicker into existence then vanish just as quickly. Moments later the dipoles reappear in a different orientation and there will be another brief dipole-dipole attraction. These momentary dipole-dipole attractions tend to be very very weak because they are only "turned on" part of the time.
The momentary dipole-dipole attractions are called instantaneous dipole-induced dipole attractions, or London forces, to distinguish them from the kind of dipole-dipole attractions that exist in polar molecules.
London forces are the only kind of attraction present between nonpolar molecules. They are also present between polar molecules and they even occur in ions. However, in polar molecules and ions these London forces contribute relatively little to the overall attractions between polar molecules and ions.
The strengths of London forces depend on several factors. One of them is the size of the electron cloud of the particle. In general, when the electron cloud is large, the outer electrons are not held very tightly by the nucleus. This makes the electron cloud "mushy" and rather easily deformed, which is precisely the condition that most favours the formation of an instanteous dipole or the creation of an induced dipole. Nuclei with large electron clouds therefore form short-lived dipoles more easily and experience stronger London forces than do similar nuclei with small electron clouds. The effects of size can be seen if you compare the boiling points of the halogens or the noble gases. Those that are large, have higher boiling points (and therefore stronger intermolecular attractions) then those that are small.
Group VIIA Boiling Point(oC) Group O Boiling Point (oC)

F2                      -188.1                   He            -268.6

Cl2                      -34.6                    Ne           -245.9

Br2                       58.8                    Ar            -185.7

I2                       184.4                    Kr            -152.3

                                                  Xe            -107.1

                                                  Rn              -61.8

A second factor that affects the strength of London forces is the number of atoms in a molecule. For molecules containing the same elements, the London forces increase with the number of atoms. The hydrocarbons are a good example. Take a look again at the alkane series and their boiling points in your databook. As the molecular mass increases by CH2 the boiling points also increase. This means that the molecule with the longer chain length experiences the stronger attractive forces. This is so because it has more places along its length where it can be attracted to other molecules. Even if the strength of attraction at each location is about the same, the total attraction experienced by the longer molecules is greater.


Hydrogen Bonds
These are a special case of dipole-dipole attraction. When hydrogen is covalently bonded to a very small, highly electronegative atom such as fluorine, oxygen, or nitrogen, unusually strong dipole-dipole attractions are often observed. There are two reasons for this. First, because of the large electronegativity difference, the F-H, O-H, or N-H bonds are very polar. The ends of these dipoles carry a substantial portion of one charge. Second, because of the small size of the atoms involved, the charge on the end of a dipole is also highly concentrated. This makes it particularly effective at attracting the opposite charge on the end of a neighbouring dipole. These two factors combine to produce attractions, called hydrogen bonds, that are about five times stronger than normal dipole-dipole attractions. Many biological molecules such as DNA and protein depend upon their shape because of these hydrogen bonds. Any factor that affects the shape of these molecules usually does so by disrupting the hydrogen bonds. Cooking an egg does this very nicely. Egg albumin is a soluble protein. Cooking it disrupts all the hydrogen bonds which reform entirely at random. The hard egg white that you eat is disrupted egg albumin protein.