IV.

Physical Properties

Physical properties of a material are those properties that do not depend on the chemical behavior of the material. Physical properties include the state of a material (gas, liquid, or solid), melting point, boiling point, crystal structure, and electrical conductivity.

A.

State

The state of a material is determined by the attraction between its atoms or molecules and by the temperature of the material. In the solid state, the attraction between the atoms or molecules is so strong that it holds them rigidly in place. The energy of vibration of the molecules of a material increases with a rise in temperature. As the temperature rises, the molecules eventually acquire enough energy to break away from their fixed positions, and the solid either melts or transforms directly into gas (a process called sublimation). The material melts if the molecular attraction remains great enough to hold the molecules together, and the material sublimes to a gas (in which the molecules are free to move randomly) if the attraction is too small.

B.

Melting Point

The melting point (or freezing point) of a substance is the temperature at which the solid form of the substance changes to a liquid (or from liquid to solid). The melting point of water is 0° on the Celsius (centigrade) temperature scale and 32° on the Fahrenheit scale (see Freezing Point).

C.

Boiling Point

The boiling point of a substance is the temperature at which the liquid form of the substance changes to a gas. The boiling point is sensitive to changes in pressure, because the molecules of a substance will tend to stay in the liquid state if they are under enough pressure. A heated liquid must overcome the atmospheric pressure in order to turn into a gas (if the atmospheric pressure exceeds the vapor pressure of the boiling liquid, the liquid will be unable to turn into vapor). For this reason, water boils at lower temperatures on high mountains (where atmospheric pressure is lower) than at sea level (where atmospheric pressure is higher). The boiling point of water at a pressure of one atmosphere, or 760 mm of mercury (a standard pressure approximating sea-level pressure), is 100° on the Celsius scale and 212° on the Fahrenheit scale.

D.

Crystal Structure

Solids may be either amorphous or crystalline in their molecular structure. In amorphous solids, the molecules are arranged haphazardly. Glass is an example of an amorphous material. Like other amorphous materials, glass does not melt at a particular temperature, because the long, randomly intertwined glass molecules cannot easily become disentangled. As a result, glass softens bit by bit as the temperature is raised, eventually becoming liquid. Crystalline materials, on the other hand, have a definite orderly array of atoms, ions, or molecules, as would a pyramid of oranges or cannonballs. The orderly arrangement of particles in a crystal is called a crystal lattice. Sand, salt, sugar, diamond, and graphite are examples of common crystalline materials. Each crystalline material has a unique melting temperature (provided the material is not chemically changed by the heat before it melts, as happens with sugar).

In an ionic crystal, the strength of mutual attraction of the ions in the crystal is reflected in the high melting point of the crystal. Table salt (or sodium chloride, NaCl), for example, melts at 800° C (1472° F). Many ionic compounds, such as sodium chloride, form crystal arrays in which each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions. The closely packed arrangement of ions in a crystalline solid, as well as the strong attraction between oppositely charged ions, accounts for the relatively hard, brittle nature of many ionic crystal solids.

Covalent crystal structures are networks of bonded atoms with the atoms occurring at the lattice points of the crystal. In the crystal lattice of diamond, each carbon atom is covalently bonded to four neighboring carbon atoms, forming a giant three-dimensional network. This three-dimensional network that composes diamond forms the hardest-known naturally occurring substance. Most covalent crystal structures are very hard and have very high melting points, because covalent bonds throughout the crystal make it essentially one giant molecule. Other examples of covalent crystals include silicon carbide (SiC), sand, and quartz (SiO₂).

Metallic crystals have unique properties because of the relationship between the positive ions and the electrons of the metal. One of the simpler and more widely used models of metallic crystals shows positive ions arranged at the points of the crystal, with electrons moving freely (as a so-called sea of electrons) among these positive ions. Because electrons in metals do not belong to any single positive ion and can move freely (carrying their electric charge with them), metals are excellent conductors of electricity. If an electric potential is applied to the metal, the electrons will move readily toward the positive electrode, creating an electric current (a stream of electrons). The freely moving electrons also make metals good transmitters of heat (metals are cold to the touch because electrons move heat away from skin).

Molecular crystals are compounds in which the molecules are held together in a crystal lattice by weak intermolecular attractive forces (for more information, see the Solutions and Solubility section of this article). These crystals do not form a complete network. Because of the weak attractions between the molecules, molecular crystals have low melting temperatures (typically well below room temperature) and are relatively soft. Most molecular organic (carbon-containing) and inorganic compounds form molecular crystals. Examples include ice (solid H₂O), solid sulfur dioxide (SO₂), and solid carbon dioxide (CO₂).