AP Chemistry - 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 Dispersion 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
Polar

HCl                -84.9                36                       18

H2S               -60.7                 34                       18

Non-polar

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. 
Crystalline Solids
The negative diagram above is of ordinary table salt on a penny. Notice that each particle is very nearly a perfect little cube. When most substances freeze, or when they separate out as a solid from a solution that is being evaporated, they normally form crystals that have highly regular features.
The particles in crystals are arranged in patterns that repeat over and over again in all directions. The overall pattern that results is called a crystal lattice. Its high degree of regularity is the principle feature that makes solids different from liquids - a liquid lacks this long range repetition of structure because the particles in a liquid are jumbled and disorganized as they move about.
Because there are millions of chemical compounds, it might seem that a enormous number of different kinds of lattices are possible. If this were true, studying solids would be hopelessly complex. Fortunately, however, the number of kinds of lattices that are mathematically possible is quite limited.
To describe the structure of a crystal it is convenient to view it as being composed of a huge number of simple, basic units called unit cells. By repeating this simple structural unit up and down, back and forth, in all directions, we can build the entire lattice. This is shown below, for the simplest and most symmetrical of all unit cells, called the simple cubic. This unit cell is a cube having atoms (or molecules or ions) at each of its eight corners. Stacking these unit cells gives a simple cubic lattice.
Two other cubic unit cells are also possible: face-centred cubic and body-centred cubic. The face-centred cubic (fcc) unit cell has identical particles at each of the corners plus another in the centre of each face.
Many common metals - copper, silver, gold, aluminum, and lead, for example - form crystals that have face-centred cubic lattices. Each of these metals has the same kind of lattice, but the sizes of their unit cells differ because the sizes of the atoms differ.
The body-centred cubic (bcc) unit cell has identical particles at each corner plus one in the centre of the cell. The body-centred cubic lattice is also common among a number of metals - examples are chromium, iron, and platinum.
Again, these are substances with the same kind of lattice, but the dimensions of the lattices reflect the size of the particular atoms.
Not all unit cells are cubic. Some have edges of different lengths or edges that intersect at angles other than 90o. Although you should realize that these other unit cells (11 other types) and the lattices they form exist, we will not go any furhter with this topic.
Below is a cutaway view of a portion of a sodium chloride crystal. The smaller particles represent Na+ ions. Notice that they are located at the lattice positions that correspond to a face-centred cubic unit cell. The Cl- ions are larger and fill the spaces between the Na+ ions. Sodium chloride is said to have a face-centred cubic lattice, and the cubic shape of this lattice is the reason that NaCl crystals take on a cubic shape when they form.
Physical Properties and Crystal Types
From personal experience you known that solids exhibit a wide range of physical properties. Some, such as diamond, are very hard; others, such as naphthalene (moth balls) or ice, are soft, by comparison, and are easily crushed. Some solids, such as salt crystals or iron, have high melting points, whereas others, such as candle wax, melt at low temperatures. Some conduct electricity well, but others are nonconducting.
 
Physical properties such as these depend on the kinds and strengths of the attractive forces that hold the particles together in the solid. Even though exact predictions cannot be made some generalizations do exist. In making these generalizations, we can divide crystals into several types according to the kinds of particles located at sites in the lattice and the kinds of attractions that exist between the particles.
 
Ionic Crystals
Ionic crystals are hard, have high melting points, and are brittle. When they melt, the resulting liquids conduct electricity well. These properties reflect the strong attractive forces between ions of opposite charge as well as the repulsions that occur when ions of like charge are placed near each other.
 
They are brittle and tend to shatter into smaller crystals when stressed. When a crystal is hammered or stressed, ions with like charges are forced into close proximity. The crystal then literaly self destructs because of electrostatic repulsion.
 
Molecular Crystals
Molecular crystals are solids in which the lattice sites are occupied either by atoms - as in solid argon or krypton - or by molecules - as in solid CO2, SO2, or H2O. Such solids tend to be soft and have low melting points because the particles in the solid experience relatively weak intermolecular attractions. The crystals are soft because little effort is needed to separate the particles or to move them past each other. The solid melts at low temperatures because the particles need little kinetic energy to break away from the solid. If the crystals contain only individual atoms, as in solid argon or krypton, or if they are composed of non-polar molecules, as in naphthalene, the only attractions between the molecules are the London forces. In crystals containing polar molecules, such as sulphur dioxide, the major forces that hold the particles together are dipole-dipole attractions. In crystals such as water the primary forces of attractions are due to hydrogen bonding.
 
Covalent or Network Crystals
Covalent crystals are solids in which the lattice points are occupied by atoms that are covalently bonded to other atoms at neighbouring lattice sites. The result is a crystal that is essentially one gigantic molecule. These solids are sometimes called network solids because of the interlocking network of covalent bonds extending throughout the crystal in all directions. A typical example is the diamond, the structure of which is illustrated below.
The above structure is diamond. Notice that each carbon atom is covalently bonded to four others at the corners of a tetrahedron. This structure extends throughout an entire diamond crystal. (In diamond, of course, all the atoms are identical. They are different shades here only to make it easier to visualize the structure.)
 
Covalent crystals tend to be hard and to have very high melting points because of the strong attractions between covalently bonded atoms. They do not conduct electricity because the electrons are bound too tightly to the bonds. Other examples of covalent crystals are quartz (SiO2 - typical grains of sand) and silicon carbide (SiC - a common abrasive used in sandpaper).
 
Metallic Crystals
Metallic crystals have properties that are quite different from those of the other three types of crystals above. They conduct heat and electricity well, and they have the lustre that we characteristically associate with metals. A number of different models have been developed to describe metals. The simplest one views the crystal as having positive ions at the lattice positions which are surrounded by electrons in a cloud that spreads throughout the entire solid.
 
The electrons in this cloud belong to no single positive ion, but rather to the crystal as a whole. Because the electrons aren't localized on any one atom, they are free to move easily, which accounts for the electrical conductivity of metals. By their movement, the electrons can also transmit kinetic energy rapidly thorough the solid, so metals are also good conductors of heat. Because the electrons are free to move easily, even the lattice points are moveable or deformable and this helps explain the malleability of metals. Even though the lattice points and electrons are free to move there are some strong attractive forces at work and these help explain ductility in some metals. This model also explains the lustre of metals. When light shines on the metal, the loosely held electrons vibrate easily and readily reemit the light with essentially the same frequency and intensity.
 
It is not possible to make simple generalizations about the melting points of metals. Some have very high melting points, like tungsten, whereas others like mercury have quite low melting points. To some degree the melting point depends on the charge of the positive ions in the metallic crystals. The ions of Group IA tend to exist as cations with a +1 charge, and they are only weakly attracted to the "electron sea" that surrounds them. Atoms of Group IIA metals, however, tend to form cations with a +2 charge. These more highly charged ions are attached more strongly to the surrounding electron sea, so the Group IIA metals have higher melting points than their neighbours in Group IA. For example magnesium melts at 650oC but sodium melts at 98oC. Tungsten which has a very high melting point, must have very strong attractions between their atoms, which suggests that there probably is some covalent bonding in them as well.
 
Crystal Types
Type 
Particles Occupying 
Lattice Sites
Type of 
Attractive Force
Typical 
Examples 
Typical 
Properties
Ionic
Positive & negative ions
Attractions between 
opposite ions 
NaCl, 
NaNO3
CaCl2
Hard; high melting points; 
non-conductors of
electricity as solids 
but good conductors when 
melted
 
Molecular
Atoms or Molecules
Dipole-dipole attractions
 London forces 
Hydrogen bonding
HCl
 SO2
N2,
Ar , 
CH4
H2
Soft, low melting  points; 
nonconductors of electricity 
in both solid and  melted form
Covalent
(network)
Atoms 
Covalent bonds 
between atoms
Diamond, 
SiC, SiO2
Very hard; very high melting pts. nonconductors of electricity
Metallic 
Atoms
Positive Attractions 
between ions and 
an electron cloud 
that extends throughout 
the crystal in both 
solid and liquid
Cu
Ag
Fe
Na
Hg
Au
Range from very hard to soft;
melting points range from high 
to low; conduct electricity and 
heat well; have a characteristic lustre
Non-Crystalline Solids
If a cubic salt crystal is broken, the pieces still have flat faces that intersect at 90o angles. On the other hand, if you shatter a piece of glass, the pieces often have surfaces that are not flat. Instead, they tend to be smooth and curved. This behaviour illustrates a major difference between crystalline solids, like NaCl, and noncrystalline solids, or amorphous solids such as glass.
The word amorphous is derived from the Greek word amorphos, which means "without form". Amorphous solids do not have long-range repetitive internal structures such as those found in crystals. In some ways their structures, being jumbled, are more like liquids than solids. Examples of amorphous solids are ordinary glass and many plastics.  In fact, glass is sometimes used as a general term to refer to any amorphous solid.
Substances that form amorphous solids usually consist of long, chainlike molecules that are intertwined in the liquid state somewhat like long strands of cooked spaghetti. To form a crystal from the melted material, these long molecules would have to become untangled and line up in specific patterns. But as the liquid is cooled, the molecules move more slowly. Unless the liquid is cooled extremely slowly, the molecular motion decreases too rapidly for the untangling to take place, and the substance solidifies with the molecules still intertwined.
Compared with substances that produce crystalline solids, those that form amorphous solids behave quite oddly when they are cooled. Those that form crystals solidify at a constant temperature. As the liquid is cooled, it eventually reaches the substances's freezing point, and crystals begin to form. Even though more heat continues to be removed, the temperature remains constant until all of the liquid has frozen. Only then does the temperature of the solid begin to drop.
Sometimes a liquid can be cooled below its freezing point. As the temperature approaches the freezing point, particles of the liquid may not be aligned in just the proper way for them to form a crystal. Thus, the temperature may continue to fall below the freezing point until, by chance, the particles in some small portion of the liquid suddenly find themselves properly arranged for a small crystal to form. This crystal grows rapidly, and the temperature rises again to the freezing point. The temperature then stays constant until all the liquid has frozen. While the liquid has a temperature below its normal freezing point, it is said to be supercooled.
With substances that form amorphous solids, the solidification of the melted material into highly ordered crystals never occurs because the molecules can't become untangled before they are frozen in place at a low temperature. Sometimes amorphous solids are therefore described as supercooled liquids. This term puts an emphasis on the kind of structural disorder found in liquids. It also suggests that the material's constituent molecules retain some residual ability at least to flex their chains into somewhat less random patterns and over a long period of time, achieve an improved degree of crystalline orderliness. But the rate at which such change occurs is typically so small that under ordinary conditions it cannot be observed and the material is fully rigid. Glass, for instance, which is a typical amorphous solid, over a very long time will develop regions of higher and higher order.
Amorphous solids also soften gradually when they are heated. This is the reason you can heat glass tubing in a flame to soften it so that it can be bent. By contrast, if you warm an ice cube (crystalline water), it won't become soft gradually - at 0OC it will suddenly melt and drip all over you!

Phase Diagrams
It is often useful to know whether a substance will be a solid, liquid or gas, at a particular temperature and pressure.  A simple way of determining this is to use a phase diagram - a graphical representation of the pressure-temperature relationships that apply to the equilibrium between the phases of the substance.
 

This
is a phase diagram for water.  As we will see, it is really not really as complicated as it appears at first glance.  On it, there are three lines that intersect at a common point. These lines given temperatures and pressures at which equilibrium between phases can exist.  The line between the brown-green phases is the vapor pressure curve for the solid (ice).  Every point on this line gives a temperature and a  pressure at which ice and its vapor are in equilibrium.  For instance, at -10oC ice has a vapor pressure of 2.15 torr.  The blue-green line represents the vapor pressure curve for liquid water.  It gives the temperatures and pressures at which the liquid and vapor are in equilibrium.  When the temperature is 100oC the vapor pressure is 760 torr.  The three lines intersect at a common point.  This is the point of temperature and pressure when only one phase can exist.  This is called the triple point.  For water the triple point occurs at 0.01oC and 4.58 torr pressure.

Every known chemical substance except helium has its own characteristic triple point, which is controlled by the balance of intermolecular forces in the solid, liquid and vapor. 

The brown-blue line which extends upwards from the triple point is the solid-liquid equilibrium line or the melting-point line.  At the triple point, melting of ice occurs at +0.01oC; at 760 torr melting occurs at 0oC.  Thus, increasing the pressure on ice lowers its melting point.
 


The decrease in the melting point of ice that occurs when there is an increse in pressure can be predicted using Le Chatelier's principle and the knowledge that liquid water is more dense then ice.  Water is very unusual as almost all other compounds have phase diagrams that show they melt at higher temperatures as the pressure increases.

The solid-liquid line in CO2 shows are marked slant to the right indicating that as the pressure increases so does the melting point temperature.

Go to the Molecular Architecture Crystals Worksheet