If the temperature of a solid is raised, according to the kinetic theory, the velocity of the particles should increase. As the temperature increases, the particles collide with each other with a greater force and, as a result, are forced farther apart. Almost all solids and liquids expand when they are heated because of this increase in velocity. If the temperature of a solid is raised sufficiently, the particles will move far enough apart to slip over one another. When such a change takes place, we say the solid has melted. The temperature at which a solid melts is called the MELTING POINT of that solid. THE NORMAL MELTING POINT IS THAT TEMPERATURE AT WHICH A SOLID BECOMES A LIQUID AT A PRESSURE OF ONE ATMOSPHERE.

The reverse is true of liquids. There will be some temperature (and pressure) at which the particles travel so slowly that they can no longer slip past one another. That temperature is called the FREEZING POINT of the liquid. THE NORMAL FREEZING POINT IS THAT TEMPERATURE AT WHICH A LIQUID BECOMES A SOLID AT 1 ATMOSPHERE OF PRESSURE. Almost all liquids have a definite freezing point, and all solids, except the amorphous forms, have a definite melting point. The freezing point of the liquid and the melting point of the solid form is the same temperature. In other words, ice melts and water freezes at the same temperature (0oC). The fact that diffusion takes place to a slight extent in the solid state and that crystalline compounds show a perceptible vapor pressure indicates that the molecules are not motionless. As the temperature of a solid is lowered, the motion of the molecules gradually decreases until at absolute zero (0oK or -273oC) it stops altogether.


Some solids have a vapor pressure great enough at room temperature that they will vaporize after a passage of time if not kept in a closed container. Such a substance will change directly from the solid to the gaseous state, without passing through the liquid state. Dry ice (solid carbon dioxide) and moth crystals (naphthalene) are substances which exhibit this phenomenon known as SUBLIMATION. It might be noted that those solids in which the intermolecular forces (van der Waals forces) are weak will exhibit measurable vapor pressures at room temperature and tend to SUBLIME. Sublimation might explain why certain solids produce characteristic odors.


A number of solids or solid-appearing substances do not possess a definite form and are referred to as being AMORPHOUS, which means "without shape". Glass is an excellent example of an amorphous solid. When glass is heated, it does not reach a point at which it suddenly becomes liquid. Glass does not have a fixed melting point as water has. Rather, as glass is heated, it softens more and more and melts gradually over a wide temperature range. The hotter it gets, the more easily the glass flows. The resistance of a liquid to flow is called its VISCOSITY. Glass and cold molasses are good examples of viscous substances. Water and carbon tetrachloride are good examples of non-viscous liquids. Upon cooling, glass does not reach a specific temperature at which it turns into a solid. As it cools, it becomes more and more viscous and it flows more and more slowly. Most amorphous substances which appear to be solids are not really solids at all. They have characteristics similar to those of glass and are called SUPERCOOLED LIQUIDS.

In order to distinguish between crystalline and amorphous substances in chemical equations, we must modify the symbol (s). In some references, true crystalline solids will be indicated by and amorphous materials by (amor).

One form of sulfur contains long chains. It is an amorphous form and is called plastic sulfur. A few hours after it is prepared, it changes back into the stable orthorhombic form. This change is characteristic of most amorphous substances. In these substances the amorphous form is unstable. Many substances, such as glass, remain in the amorphous form for long periods instead of changing to a more stable crystalline form as sulfur does. Substances which can occur in long-lasting amorphous form are said to be METASTABLE. Although a substance in the metastable form is not in its most stable form, it is not likely to change unless subjected to some external disturbance. Glass is a metastable substance which normally occurs in the amorphous form, but even glass may be crystallized under proper conditions.


Most solid substances are crystalline in nature. Some solids, such as sugar and table salt in the form in which we are most familiar, are composed of single crystals. The majority of crystalline solids with which we come in daily contact are aggregates of many interlocking small crystals. A CRYSTAL MAY BE DEFINED AS A HOMOGENEOUS BODY HAVING THE NATURAL SHAPE OF A POLYHEDRON. Crystals are three-dimensional figures bounded by plane surfaces. The angles at which the surfaces intersect are always the same for a given substance, and are characteristic of that substance. This statement is often called STENO'S LAW named after Nicolaus Steno who first made the observation in 1669. The structure of crystalline solids was made clear in 1912 when a German chemist named Max von Laue discovered that the arrangement of the atoms, ions, or molecules in crystals could be determined by means of X-rays. Crystals of each substance are built up according to a fundamental geometric pattern known as the CRYSTAL LATTICE. The smallest fraction of the crystal lattice which contains a representative portion of the crystal structure is referred to as a UNIT CELL. The nature of the solid is determined therefore, by the size, shape, and content of its unit cell.

The unit cells for known crystals can be grouped into seven types:

Unit Cell Example
Cubic Rock Salt
Tetragonal White Tin
Orthorhombic Mercury (II) Chloride
Monoclinic Potassium Chlorate
Triclinic Potassium Dichromate
Hexagonal Silica
Rhombohedral Calcite

Variations of the seven types of unit cells give rise to fourteen observed crystal lattices:

Cubic, Simple Orthorhombic, base-centered
Cubic, body-centered Orthorhombic, body-centered
Cubic, face-centered Orthorhombic, face-centered
Triclinic Hexagonal
Monoclinic Rhombohedral
Monoclinic Tetragonal, simple
Orthorhombic, simple Tetragonal, body-centered

A crystal lattice is said to be BODY-CENTERED if each unit cell contains at its center an atom of the same kind as at the corners. The lattice is FACE-CENTERED if the unit cell contains at the midpoints of its faces atoms of the same kind as at the corners. A crystal lattice is said to be BASE-CENTERED if each unit cell contains at the midpoint of its most extreme faces atoms of the same kind as at the corners. The structure of a crystal is determined by various factors. Among these are the relative numbers and sizes of the building units and the types of bonds holding the building units together in the crystal. The building units may be atoms, ions, or molecules. Copper forms atomic crystals which are built up on copper atoms packed together in a regular pattern (face-centered cubic) with the atoms occupying the lattice positions. Solid carbon dioxide (dry ice) is composed of molecular crystals in which CO2 molecules are the unit particles. Sodium chloride forms ionic crystals in which sodium ions and chloride ions are the building units.

Crystals of NaF, KCl, MgO, and CaS all have the same crystal structure as NaCl. SrCl2, PbF2, and CdF2 all crystallize with the structure characteristic of CaF2. Different compounds which crystallize with the same structure are said to be ISOMORPHOUS. When calcium carbonate (CaCO3) crystallizes at low temperatures it assumes a Rhombohedral structure and is called calcite, but when it crystallizes at high temperatures an orthorhombic lattice results and the substance is called aragonite. The ability to assume two or more crystalline structures by the same substance is called POLYMORPHISM. It should be remembered that some of the atoms shown in a diagram of the unit cell are actually shared by other unit cells and do not lie in their entirety within one unit cell. The following rules may help:

  1. An atom which lies completely within the unit cell belongs to that cell only.
  2. An atom lying on a face of a unit cell belongs equally to two unit cells and counts as one-half of one atom for a unit cell.
  3. An atom lying on an edge is shared equally by four unit cells and counts as one-fourth of an atom for a unit cell.
  4. An atom lying at a corner is shared equally by eight unit cells, except in the hexagonal cell, and counts as one-eighth on an atom for each unit cell.
  5. An atom lying at a corner of a hexagonal cell is shared equally by six unit cells and counts as one-sixth of an atom for each unit cell.


When the temperature of crystalline solid reaches the melting point, it remains constant until all of the solid has changed to the liquid state. The quantity of heat required to change a unit mass of a substance from the solid to the liquid state at constant temperature is known as the HEAT OF FUSION. The heat of fusion of ice is 80 calories per gram (1440 cal/mole). Different substances have different heats of fusion just as different substances have different heats of vaporization (Section 7.4).

Table: Heat of Fusion and Vaporization
Substance Melting Point (oC) Boiling Point (oC) Solid <--> Liquid Heat of Fusion Liquid <-->Vapor Heat of Vaporization
Oxygen -218 -183 3.3 cal/g 51 cal/g
Nitrogen -210 -196 6.1 cal/g 48 cal/g
Alcohol -114 78 26.0 cal/g 204 cal/g
Water 0 100 80.0 cal/g 540 cal/g
Lead 327 1750 5.5 cal/g 205 cal/g
Silver 961 2212 26.5 cal/g 563 cal/g


Some of the general properties of solids which we can easily see are:

  1. DEFINITE SHAPE. A solid maintains its shape. Unlike liquids and gases, it does not flow under ordinary conditions. The shape of a solid is independent of its container.
  2. DEFINITE VOLUME. All of the surfaces of a solid are free surfaces. Thus, the volume of a solid is also independent of its container.
  3. NONCOMPRESSIBILITY. The pressures required to decrease the volumes of solids are even greater than those required for liquids.
  4. VERY SLOW DIFFUSION. If a lead plate and a gold plate are placed in close contact particles of gold can be detected in the lead and vice versa.
  5. CRYSTAL FORMATION. Solids may be described as either crystalline or amorphous. Crystalline solids have a regular arrangement of particles. Amorphous solids have a completely random arrangement of particles.


It is frequently more useful to classify crystals according to the type of lattice structure bonding that they possess.

  1. IONIC CRYSTAL. The ionic crystal lattice consists of positive and negative ions arranged in a characteristic regular pattern. No molecular units are evident with-in the crystal. The strong bonding forces result from the attraction of positive and negative charges. Ionic crystals are hard and brittle, have rather high melting points, and are good insulators.
  2. COVALENT NETWORK CRYSTAL. The covalent network crystal lattice consists of an array of atoms that share electrons with neighboring atoms. The binding forces are strong covalent bonds which extend in fixed directions. The resulting crystals are compact, interlocking, covalent network structures. They may be considered to be giant molecules. They are very hard and brittle, have rather high melting points, and are nonconductors. Diamond, silicon carbide, silicon dioxide, and oxides of transition metals are of this type.
  3. METALLIC CRYSTAL. The metallic crystal lattice consists of positive ions surrounded by a cloud of valence electrons. This cloud is commonly referred to as the electron gas. The binding force is the attraction between the positive ions of the metal and the electron "gas". The valence electrons are donated by the atoms of the metal and belong to the crystal as a whole. These electrons are free to migrate throughout the crystal lattice. This electron mobility explains the high electric conductivity associated with metals. The hardness and melting points of metallic crystals vary greatly for different metals. Sodium, iron, tungsten, copper, and silver are typical examples of metallic crystals that have good electric conductivity.
  4. COVALENT MOLECULAR CRYSTAL. The covalent molecular crystal lattice consists of an orderly arrangement of individually distinct molecules. If the molecules are Nonpolar, the binding force is relatively weak. If the molecules are polar, the binding force is relatively strong. The covalent chemical bonds which bind the atoms within the molecules are much stronger than the forces which form the crystal lattice. Therefore, molecular crystals have low melting points, are relatively soft, volatile (easily vaporized), and good insulators. Iodine, carbon dioxide, water, and hydrogen form crystals of this type.


A perfect crystal is a rarity. Most crystals contain defects of one or more types. If we look at a plane of a simple crystal such as sodium chloride, one of the ions may be removed from its proper position and occupy a space where no ion usually occurs. This causes an imperfect crystal. Another possibility is that an ion is missing completely from its position in the lattice. In the event that a defect of this type occurs, for every positive ion missing there must be a negative ion missing, in order to preserve the electrical neutrality of the crystal. It is sometimes possible for extraneous ions, atoms, electrons, or even molecules to occupy these spaces vacated by the normal ions of the crystal.

Controlled minute amounts of impurities are sometimes deliberately introduced into the crystal to cause imperfections in the lattice. These impurity defects result in changes in the conductivity of the crystal which gives rise to practical applications such as in the manufacture of semiconductors.


When aqueous solutions of many soluble salts are evaporated, the salt separates as crystals which contain the salt and water combined in definite proportions by weight. Such compounds are known as HYDRATES and the water is called WATER OF HYDRATION. The formation of hydrates is not limited to salts but is common with acids and bases and even elements. In fact it is not confined to crystals. Some examples of hydrates are:

CuSO4 . 5 H2O Blue vitriol
MgSO4 . 7 H2O Epsom salts
KAl(SO4)2 . 12 H2O Alum
Na2SO4 . 10 H2O Glauber's salt

When blue crystals of copper (II) sulfate-5-hydrate, CuSO4 . 5 H2O, are heated, water is produced, the crystalline structure characteristic of the hydrated salt breaks down, and the white anhydrous (without water) salt, CuSO4, remains.

CuSO4 . 5 H2O rarrow.gif (63 bytes) CuSO4 + 5 H2O


Hydrates exhibit vapor pressure which can be demonstrated by placing a crystal of a hydrate in the vacuum above the mercury in a barometric tube, the mercury level being depressed to an extent dependent upon the magnitude of the VAPOR PRESSURE OF THE HYDRATE. The vapor pressure of a hydrate increases with an increase in temperature and decreases as the temperature is lowered.

When a hydrate exhibits a vapor pressure higher than the partial pressure of water in the atmosphere, it will loose a part or all of its water of hydration when exposed to the air. Hydrates which loose water of hydration when exposed to the air are said to be EFFLORESCENT, and EFFLORESCENCE is said to have taken place. It follows that some hydrates are stable when the humidity is high but decompose when it is low. A salt such as copper (II) sulfate may form more than one hydrate, each of which possesses its own definite vapor pressure at a given temperature. The following hydrates of copper (II) sulfate are known:

CuSO4 . 5 H2O CuSO4 . 3 H2O CuSO4 . H2O

  When certain substances of low vapor pressure, such as CaCl2 . H2O, are exposed to air, they form higher hydrates. Such salts may be used in the removal of moisture from air or other gases. A substance that can remove moisture from the air is said to be HYGROSCOPIC. Concentrated sulfuric acid, a liquid, and phosphorus (V) oxide, a solid, are powerful drying agents, or DESICCANTS. Certain water-soluble hygroscopic solids remove sufficient water from the air to dissolve completely in this water and form solution (P2O5). Such substances are said to be DELIQUESCENT, and the process is termed DELIQUESCENCE. Very soluble salts, such as CaCl2 are often extremely deliquescent.

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Copyright May 1987 James R. Fromm (mailto:jfromm@3rd1000.com) Revised February 2000