All atoms in covalent network solids are covalently bonded. The high melting points, hardness, and other properties are explained by covalently-bonded atoms forming three, two, and one dimensional networks, e.g., quartz, mica, and asbestos.

Many properties of molecular solids and liquids are due to van der Waals forces. These are relatively weak forces between neighboring molecules either due to momentary dipoles or due to electrostatic interactions among polar molecules. Van der Waals forces are commonly classified as either London forces (between molecules of hydrogen and carbon dioxide for example) or dipole-dipole interactions (between molecules of hydrogen chloride and chloroform for example). Hydrogen bonding, a special case of dipole-dipole interaction, occurs between hydrogen atoms bonded to nitrogen, oxygen, or fluorine atoms in one molecule and a nonbonding electron pair on the nitrogen, oxygen, or fluorine atom in an adjacent molecule (as in water and ammonia, for example). In some cases hydrogen bonds can form within the same molecule, depending on its geometry and composition.

Place in the Curriculum

Bonding logically follows development of a detailed atomic model of matter including electron configurations and a consideration of chemical periodicity. The topic of bonding may be treated as a whole or may be partitioned into several segments. Personal preference or the organization of the adopted textbook may dictate order of coverage. In any case, ionic and covalent bonding should be treated, while covalent networks, van der Waals forces, and metallic bonding may be deferred until condensed phases are taught.

Central Concepts

1. The electrons in the highest occupied energy level in a ground-state atom determine the atom's chemical properties and are called valence electrons.

2. Atoms may form molecules or aggregates and gain stability by sharing an electron pair, giving rise to an electrostatic attraction, thus forming a covalent bond.

3. Metals and nonmetals may form a stable structure by transferring electrons from the metal atom to the nonmetal atom. This transfer forms oppositely-charged particles called ions; the resulting structure is held together by ionic bonds.

4. Valence electrons are arranged around the kernels (the nucleus and electrons other than the valence electrons) of atoms, forming molecules. Valence electrons can be represented as dots, yielding structures (Lewis-dot structures) consistent with the octet rule.

5. Carbon, nitrogen, oxygen, and sulfur atoms frequently share two or three electron pairs, forming double or triple covalent bonds consistent with the octet rule.

6. Like atoms (i.e., same electronegativity) form "equal-sharing" covalent bonds, called nonpolar bonds. Atoms with different electronegativities do not equally share electron pairs; they form polar covalent bonds.

7. When two isolated atoms capable of bonding approach each other, the potential energy of the system decreases due to attraction and is a minimum at some equilibrium distance (called the bond distance) separating the atoms. If the atoms attempt to approach more closely, the attraction changes to repulsion rapidly, hence the potential energy of the system also rises rapidly.
8. Positively-charged atomic kernels in metals are held in closely-packed crystal structures by mobile valence electrons.
9. Covalent network solids involve atoms covalently bonded in three, two, or one dimensional networks. The primary covalent bonds give these substances many of their properties.
10. Intermolecular (as opposed to intramolecular) attractions are responsible for aggregates of molecules which become solids or liquids. The weakest attractions are collectively called van der Waals forces. London forces, the weakest of these forces, are caused by momentary fluctuations of electron distribution symmetry in the atoms. The strength of theses forces depends on polarizability of the electron distributions and surface area of the molecules. Although London forces are the weakest of the van der Waals forces, they may be considered the most important. This is because they occur between every atom or molecule regardless of whether or not any other forces are present. Dipole-dipole interactions occur when polar molecules are attracted to each other. The strongest dipole-dipole interactions occur when a hydrogen atom, covalently bonded to nitrogen, oxygen, or fluorine, is attracted to an unshared electron pair of a similar highly electronegative atom in the same or adjacent molecule. This is called hydrogen bonding.

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