As seen in the many examples describing the properties of solids, the particles that make up a solid and the types of bonds with which they are held together are critical pieces of information needed in predicting the relative behavior of a material. Accordingly, solids can be categorized based on those characteristics. Table 7.6 below summarizes the different types of solids and the properties generally associated with them.
Table 7.6 Types of Solids
| Type of Solid | Examples | Pertinent Particles |
Type of Bonds | General Properties |
| Macromolecular (Also known as a network solid) | C (as in diamond or graphite), SiC, SiO2 | Nonmetal atoms | Covalent | High heat of fusion and melting point; low electrical conductivity and solubility in water |
| Ionic | NaCl, MgSO4, K2Cr2O7 | Metal and nonmetal ions | Ionic | High heat of fusion, melting point, and solubility in water*; low electrical conductivity |
| Metallic | Cu, Mg, Hg | Metal atoms | Metallic | Moderate to high heat of fusion and melting point; high electrical conductivity; low solubility in most solvents |
| Molecular | SO2, NH3, CH4 | Molecules | Intermolecular | Low heat of fusion, melting point, and electrical conductivity; variable solubility in water |
*There are many ionic substances insoluble in water.
There is a special type of solid that contains ions and molecules combined in a single substance; it is called a hydrate. Hydrates can be viewed as containing fully charged ions and polar water molecules surrounded by and attracted to each other in a crystal structure that appears dry despite the presence of the water. A hydrate’s formula displays the formula for the regular ionic substance followed by the number of water molecules associated with it per formula unit. Examples include CuSO4 • 5H2O and BaCl2 • 2H2O (the • in the formula is read as “with”). Often, the heating of hydrated crystals releases their water of hydration and alters the crystal structure of the remaining “anhydrous” ionic material; this in turn causes a color change in the solid substance. For example, if a sample of CuSO4 • 5H2O (called copper (II) sulfate pentahydrate) is heated in the bottom of a test tube, the blue solid turns into a white powder with liquid water seen accumulating near the cooler top of the test tube. Certain hydrated substances—for example, MgSO4 • 7H2O (Epsom salt)—lose their water of hydration at room temperature without the need for additional heating in a process called efflorescence. Conversely, other ionic solids are so attracted to water molecules that they readily become hydrates by absorbing moisture from the air; these are referred to as hygroscopic materials. Calcium chloride is such a substance and is commonly used as a drying agent in the laboratory because of its desire to be associated with water. Some ionic substances are so hygroscopic that they can become wet and form “puddles” of ionic solutions with the water they accumulate. These ionic substances are additionally referred to as being deliquescent. Sodium hydroxide is an example of such a substance.