The physical properties of crystalline solids, such as melting point and hardness, depend both on the arrangements of particles and on the attractive forces between them. Table 11.6 classifies solids according to the types of forces between particles in solids.
Molecular solids consist of atoms or molecules held together by intermolecular forces (dipole-dipole forces, London dispersion forces, and hydrogen bonds). Because these forces are weak, molecular solids are soft. Furthermore, they normally have relatively low melting points (usually below 200°C). Most substances that are gases or liquids at room temperature form molecular solids at low temperature. Examples include Ar, H2O, and CO2.
The properties of molecular solids depend not only on the strengths of the forces that operate between molecules but also on the abilities of the molecules to pack efficiently in three dimensions. For example, benzene, C6H6, is a highly symmetrical planar molecule. It has a higher melting point than toluene, a compound in which one of the hydrogen atoms of benzene has been replaced by a CH3 group (Figure 11.39). The lower symmetry of toluene molecules prevents them from packing as efficiently as benzene molecules. As a result, the intermolecular forces that depend on close contact are not as effective, and the melting point is lower. In contrast, the boiling point of toluene is higher than that of benzene, indicating that the intermolecular attractive forces are larger in liquid toluene than in liquid benzene. Figure 11.39 also presents the melting and boiling points of another substituted benzene phenol. Both the melting and boiling points of phenol are higher than those of benzene because of the hydrogen-bonding ability of the OH group of phenol.
Figure 11.39 Comparative melting and boiling points for benzene, toluene, and phenol.
Covalent-network solids consist of atoms held together in large networks or chains by covalent bonds. Because covalent bonds are much stronger than intermolecular forces, these solids are much harder and have higher melting points than molecular solids. Two of the most familiar examples of covalent-network solids are diamond and graphite, two allotropes of carbon. Other examples include quartz, SiO2, silicon carbide, SiC, and boron nitride, BN.
In diamond each carbon atom is bonded to four other carbon atoms as shown in Figure 11.40(a). This interconnected three-dimensional array of strong carbon-carbon single bonds contributes to diamond's unusual hardness. Industrial-grade diamonds are employed in the blades of saws for the most demanding cutting jobs. In keeping with its structure and bonding, diamond also has a high melting point, 3550°C.
Figure 11.40 Structures of (a) diamond and (b) graphite.
In graphite the carbon atoms are arranged in layers of interconnected hexagonal rings as shown in Figure 11.40(b). Each carbon atom is bonded to three others in the layer. The distance between adjacent carbon atoms in the plane, 1.42 Å, is very close to the C C distance in benzene, 1.395 Å. In fact, the bonding resembles that of benzene, with delocalized bonds extending over the layers. Electrons move freely through the delocalized orbitals, making graphite a good conductor of electricity along the layers. (If you have ever taken apart a flashlight battery, you know that the central electrode in the battery is made of graphite.) The layers, which are separated by 3.41 Å, are held together by weak dispersion forces. The layers readily slide past one another when rubbed, giving the substance a greasy feel. Graphite is used as a lubricant and in making the "lead" in pencils. Figure 11.41 is a photomicrograph of graphite that reveals its layered structure.
Ionic solids consist of ions held together by ionic bonds. The strength of an ionic bond depends greatly on the charges of the ions. Thus, NaCl, in which the ions have charges of 1+ and 1-, has a melting point of 801°C, whereas MgO, in which the charges are 2+ and 2-, melts at 2852°C.
The structures of simple ionic solids can be classified as a few basic types. The NaCl structure is a representative example of one type. Other compounds that possess this same structure include LiF, KCl, AgCl, and CaO. Three other common types of crystal structures are shown in Figure 11.42.
Figure 11.42 Unit cells of some common types of crystal structures found for ionic solids: (a) CsCl; (b) ZnS (zinc blende); (c) CaF2 (fluorite).
The structure adopted by an ionic solid depends largely on the charges and relative sizes of the ions. In the NaCl structure, for example, the Na+ ions have a coordination number of 6: Each Na+ ion is surrounded by six nearest neighbor Cl– ions. In the CsCl structure [Figure 11.42(a)], by comparison, the Cl– ions adopt a simple cubic arrangement with each Cs+ ion surrounded by 8 Cl– ions. The increase in the coordination number as the alkali metal ion is changed from Na+ to Cs+ is a consequence of the larger size of Cs+ compared to Na+.
In the zinc blende, ZnS, structure [Figure 11.42(b)], the S2– ions adopt a face-centered cubic arrangement, with the smaller Zn2+ ions arranged so they are each surrounded tetrahedrally by four S2– ions. CuCl also adopts this structure.
In the fluorite, CaF2, structure [Figure 11.42(c)] the Ca2+ ions are shown in a face-centered cubic arrangement. As required by the chemical formula of the substance, there are twice as many F– ions in the unit cell as there are Ca2+ ions. Other compounds that have the fluorite structure include BaCl2 and PbF2.
Metallic solids consist entirely of metal atoms. Metallic solids usually have hexagonal close-packed, cubic close-packed (face-centered cubic), or body-centered cubic structures. Thus, each atom typically has 8 or 12 adjacent atoms.
The bonding in metals is too strong to be due to London dispersion forces, and yet there are not enough valence electrons for ordinary covalent bonds between atoms. The bonding is due to valence electrons that are delocalized throughout the entire solid. In fact, we can visualize the metal as an array of positive ions immersed in a sea of delocalized valence electrons, as shown in Figure 11.45.
Figure 11.45 A cross section of a metal. Each sphere represents the nucleus and inner-core electrons of a metal atom. The surrounding colored "fog" represents the mobile sea of electrons that binds the atoms together.
Metals vary greatly in the strength of their bonding, as is evidenced by their wide range of physical properties such as hardness and melting point. In general, however, the strength of the bonding increases as the number of electrons available for bonding increases. Thus, sodium, which has only one valence electron per atom, melts at 97.5°C, whereas chromium, with six electrons beyond the noble-gas core, melts at 1890°C. The mobility of the electrons explains why metals are good conductors of heat and electricity. The bonding and properties of metals will be examined more closely in Chapter 23.
A normal breath of air has a volume of about 500 mL. If the water vapor in the air is at the equilibrium vapor pressure for the water, we say that the air is saturated with water vapor. How many molecules of water are present in a breath of 508 mL if it is saturated with water at normal body temperature, 37°C? (The vapor pressure of water at different temperatures is given in Appendix B of the textbook.)
SOLUTION The number of molecules can be determined from the number of moles of water vapor in the breath. The number of moles can be calculated using the ideal gas equation, PV = nRT.
We are given the volume of the gas, V = 508 mL = 0.508 L and the temperature, T = 37 + 273 = 310 K, and we are informed that we can obtain the partial pressure of the water vapor from Appendix B. In Appendix B we find that the vapor pressure of water is 42.2 torr at 35°C and 55.3 torr at 40°C. We can estimate the vapor pressure at 37°C by interpolating between the given values. We estimate the vapor pressure to be of the way between the 35°C and 40°C values just as 37°C is the way between these temperatures. The difference between the vapor pressures at the two tabulated temperatures is 55.3 torr - 42.2 torr = 13.1 torr. Thus, we estimate the vapor pressure at 37°C to be
Solving the ideal gas equation for nH2O gives:
Using Avogadro's number, we then convert moles to molecules.