Network atomic solids forces

The atoms in network atomic solids are held together by forces created when electrons are shared between atoms. These forces create a type of chemical bond known as a covalent bond. 

The atoms in molecular solids are held together by weak inter-molecular forces. These forces are much weaker than the chemical bonds in ionic, metallic, and network atomic solids, but they are still strong enough to hold molecules together. 

There can be slightly different forces holding particles together within a solid. Ionic solids, metallic solids, network atomic solids, molecular solids, and amorphous solids each use a different force or combination of forces to hold molecules or atoms together. 

Crystalline solids can be classified into five categories based on the types of particles they contain atomic solids, molecular solids, covalent network solids, ionic solids, and metallic solids. Table 13-4 summarizes the general characteristics of each category and provides examples. The only atomic solids are noble gases. Their properties reflect the weak dispersion forces between the atoms. 

The type of attractive forces within solids depends on the identity of the unit particle and the chemical bonds it can form. The forces between atoms in a covalent network solid (such as carbon in diamond) are covalent bonds. These bonds result when at least one pair of electrons is shared by two atoms. The forces between atoms within metallic elements (such as iron) are metallic bonds. Electrostatic attractions—also called ionic bonds—are the forces between ions, atoms which have lost one or more electrons to become positively charged ions or which have gained one or more electrons to become negatively charged ions (such as those found in NaCI). Ionic compounds are often known as salts. Covalent, metallic, and ionic bonds are strong chemical bonds. 

Molecular covalent substances are soft and low melting because of the weak forces between the molecules. Network covalent solids are hard and high melting because covalent bonds join all the atoms in the sample. 

Solids whose composite units are individual atoms are atomic solids. SoUd xenon (Xe), iron (Fe), and silicon dioxide (Si02) are examples of atomic soUds. We can classify atomic solids themselves into three categories—nonbondmg atomic solids, metallic atomic solids, and network covalent atomic solids—each held together by a different kind of force. 

Network covalent atomic solids, such as diamond, graphite, and silicon dioxide, are held together by covalent bonds. The crystal structures of these solids are more restricted by the geometrical constraints of the covalent bonds (which tend to be more directional than intermolecular forces, ionic bonds, or metallic bonds) so they do not tend to form closest-packed structures. 

We have seen that the pure elements may solidify in the form of molecular solids, network solids, or metals. Compounds also may condense to molecular solids, network solids, or metallic solids. In addition, there is a new effect that does not occur with the pure elements. In a pure element the ionization energies of all atoms are identical and electrons are shared equally. In compounds, where the most stable electron distribution need not involve equal sharing, electric dipoles may result. Since two bonded atoms may have different ionization energies, the electrons may spend more time near one of the positive nuclei than near the other. This charge separation may give rise to strong intermolecular forces of a type not found in the pure elements. 

The molecules (or atoms, for noble gases) of a molecular solid are held In place by the types of forces already discussed In this chapter dispersion forces, dipolar interactions, and/or hydrogen bonds. The atoms of a metallic solid are held in place by the delocalized bonding described in Section 10-. A network solid contains an array of covalent bonds linking every atom to its neighbors. An ionic solid contains cations and anions, attracted to one another by electrical forces as described in Section 8-. 

Many of the solids around us are ionic. They include most of the minerals that form our landscapes. Many others are molecular, in which electrically neutral molecules stack together under the influence of intermole-cular forces, as in ice and glucose. Solids also include network solids such as diamond and quartz, in which all the atoms in an entire crystalline region are covalently bonded in place. 

The big difference in melting points suggests a difference in type of crystal binding. The intermolecular forces in solid CO2 must be very low to be overcome by a low-temperature sublimation. CO2 is actually a molecular lattice held together only by the weak van der Waals forces between discrete CO2 molecules. Si02 is a covalent lattice with a three-dimensional network of bonds each silicon atom is bonded tetrahedrally to four oxygen atoms and each oxygen is bonded to two silicon atoms. 

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