10.7 Structure, Properties and Bonding in Solids

1. Different Types of Solids

• Metallic solids are held together by a “sea” of collectively shared electrons.
• Ionic solids are sets of cations and anions mutually attracted to one another.
• Covalent-network solids are joined by an extensive network of covalent bonds.
• Molecular solids are discrete molecules held together by weak forces.

Two Other Types of Solids

Polymers contain long chains of atoms connected by covalent bonds; the chains can be connected to other chains by weak forces.  These molecules have different properties than small molecules or metallic or ionic compounds.

Nanomaterials are crystalline compounds with the crystals on the order of 1–100 nm; this gives them very different properties than larger crystalline materials.

2. Organization of Solids

• Solids with a regular repeating pattern of atoms are crystalline.
• Amorphous solids are characterized by a distinct lack of order in the arrangement of atoms.

Since crystalline solids have a regular pattern, they are of more interest to most chemists.

Figure 10.25

Source: commons.wikimedia.org/

3. The solid state of matter

The solid state, being the form of any substance that prevails at lower temperatures, is one in which thermal motion plays an even smaller role than in liquids. The thermal kinetic energy that the individual molecular units do have at temperatures below their melting points allows them to oscillate around a fixed center whose location is determined by the balance between local forces of attraction and repulsion due to neighboring units, but only very rarely will a molecule jump out of the fixed space alloted to it in the lattice. Thus solids, unlike liquids, usually exhibit long-range order, cohesiveness and rigidity, and possess definite shapes.

Classification of matter according to bond type

The most commonly used classification is based on the kinds of forces that join the molecular units of a solid together. We can usually distinguish four major categories on the basis of properties such as general appearance, hardness, and melting point.

Table 10.8

Figure 10.26

Classification by type of molecular unit

Solids, like the other states of matter, can be classified according to whether their fundamental molecular units are atoms, electrically-neutral molecules, or ions. But solids possess an additional property that gases and liquids do not: an enduring structural arrangement of their molecular units. Classification by dominant attractive force

Notice how the boiling points in the following selected examples reflect the major type of attractive force that binds the molecular units together. Bear in mind, however, that more than one type of attractive force can be operative in many substances. This topic is discussed in more detail here.

Table 10.9

4. The nature of crystalline solids

In a solid comprised of identical molecular units, the most favored (lowest potential energy) locations occur at regular intervals in space. If each of these locations is actually occupied, the solid is known as a perfect crystal.

Much more on crystals and their structures can be found in these units that deal with lattices and external shapes, cubic crystals and packing of spheres, ionic solids, and determination of crystal structures.

What really defines a crystalline solid is that its structure is composed of repeating unit cells each containing a small number of molecular units bearing a fixed geometric relation to one another. The resulting long-range order defines a three-dimensional geometric framework known as a lattice. [image]

Figure 10.27

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Geometric theory shows that only fourteen different types of lattices are possible in three dimensions, and that just six different unit cell arrangements can generate these lattices. The regularity of the external faces of crystals, which in fact correspond to lattice planes, reflects the long-range order inherent in the underlying structure.

Perfection is no more attainable in a crystal than in anything else; real crystals contain defects of various kinds, such as lattice positions that are either vacant or occupied by impurities, or by abrupt displacements or dislocations of the lattice structure.

Most pure substances, including the metallic elements, form crystalline solids. But there are some important exceptions.

Metallic solids

The details of metallic bonding are covered in this Chemical Bonding unit.

In metals the valence electrons are free to wander throughout the solid, instead of being localized on one atom and shared with a neighboring one. The valence electrons behave very much like a mobile fluid in which the fixed lattice of atoms is immersed. This provides the ultimate in electron sharing, and creates a very strong binding effect in solids composed of elements that have the requisite number of electrons in their valence shells. The characteristic physical properties of metals such as their ability to bend and deform without breaking, their high thermal and electrical conductivities and their metallic sheen are all due to the fluid-like behavior of the valence electrons.

Molecular solids

Recall that a “molecule” is defined as a discrete aggregate of atoms bound together sufficiently tightly (that is, by directed covalent forces) to allow it to retain its individuality when the substance is dissolved, melted, or vaporized.

The two words italicized in the preceding sentence are important; covalent bonding implies that the forces acting between atoms within the molecule are much stronger than those acting between molecules, and the directional property of covalent bonding confers on each molecule a distinctive shape which affects a number of its properties. Most compounds of carbon — and therefore, most chemical substances, fall into this category.

Many simpler compounds also form molecules; H₂O, NH₃, CO₂, and PCl₅ are familiar examples. Some of the elements, such as H₂, O₂, O₃, P₄ and S₈ also occur as discrete molecules. Liquids and solids that are composed of molecules are held together by van der Waals forces, and many of their properties reflect this weak binding. Thus molecular solids tend to be soft or deformable, have low melting points, and are often sufficiently volatile to evaporate (sublime) directly into the gas phase; the latter property often gives such solids a distinctive odor.

Figure 10.28

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Iodine is a good example of a volatile molecular crystal. The solid (mp 114° C , bp 184°) consists of I₂ molecules bound together only by dispersion forces. If you have ever worked with solid iodine in the laboratory, you will probably recall the smell and sight of its purple vapor which is easily seen in a

closed container.

Figure 10.29

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Because dispersion forces and the other van der Waals forces increase with the number of atoms, larger molecules are generally less volatile, and have higher melting points, than do the smaller ones. Also, as one moves down a column in the periodic table, the outer electrons are more loosely bound to the nucleus, increasing the polarisability of the atom and thus its susceptibility to van der Waals-type interactions. This effect is particularly apparent in the progression of the boiling points of the successively heavier noble gas elements.

Covalent solids

Figure 10.30

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The diamond lattice structure

There is a class of extended-lattice compounds (see Section 6 below) in which each atom is covalently bonded to its nearest neighbors. This means that the entire crystal is in effect one super-giant “molecule”.

The diamond form of elemental carbon is the most well-known example of a covalent solid. Each carbon atom is directly bonded to four adjacent atoms.

This tight bonding accounts for the extreme hardness of these solids; they cannot be broken or abraded without cleaving a large number of covalent chemical bonds. Similarly, a covalent solid cannot “melt” in the usual sense; when heated to very high temperatures, these solids usually decompose into their elements.

More about diamond

Diamond is the hardest material known, defining the upper end of the 1-10 scale known as Moh’s hardness. Diamond cannot be melted; above 1700°C it is converted to graphite, the more stable form of carbon.

Figure 10.31

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The diamond unit cell is face-centered cubic and contains 8 carbon atoms. The four darkly shaded ones are contained within the cell and are completely bonded to other members of the cell. The other carbon atoms (6 in faces and 4 at corners) have some bonds that extend to atoms in other cells. (Two of the carbons nearest the viewer are shown as open circles in order to more clearly reveal the bonding arrangement.)

Other covalent solids

Boron nitride BN is similar to carbon in that it exists as a diamond-like cubic polymorph as well as in a hexagonal form analogous to graphite. Cubic BN is the second hardest material after diamond, and finds use in industrial abrasives and cutting tools. Recent interest in BN has centered on its carbon-like ability to form nanotubes and related nanostructures. (A-Zmaterials article, Nanotechnology article, Wikipedia article)

Silicon carbide SiC is an extremely rare mineral on the earth, and comes mostly from meteorites which are believed to have their origins in carbonaceous stars. The first synthetic SiC was made accidently by E.G. Acheson in 1891 who immediately recognized its industrial prospects and founded the Carborundum Co.

A commercial grade of silicon carbode is known as carborundum. Its structure is very much like that of diamond with every other carbon replaced by silicon. On heating at atmospheric pressure, it decomposes at 2700°C, but has never been observed to melt. Structurally, it is very complex; at least 70 crystalline forms have been identified. Its extreme hardness and ease of synthesis have led to a diversity of applications — in cutting tools and abrasives, high-temperature semiconductors, and other high-temperature applications, manufacture of specialty steels, jewelry, and many more.

Tungsten carbide WC is probably the most widely-encountered covalent solid owing to its use in “carbide” cutting tools and as the material used to make the rotating balls in ball-point pens. It’s high-melting (2870°C) form has a structure similar to that of diamond and is only slightly less hard. In many of its applications it is embedded in a softer matrix of cobalt or coated with titanium compounds. (Wikipedia article)

5. Amorphous solids

In some solids there is so little long-range order that the substance cannot be considered crystalline at all; such a solid is said to be amorphous. Amorphous solids possess short-range order but are devoid of any organized structure over longer distances; in this respect they resemble liquids. However, their rigidity and cohesiveness allow them to retain a definite shape, so for most practical purposes they can be considered to be solids.

Glasses

This term refers generally to solids formed from their melts that do not return to their crystalline forms on cooling, but instead form hard, and often transparent amorphous solids. Although some organic substances such as sugar can form glasses (“rock candy”), the term more commonly describes inorganic compounds, especially those based on silica, SiO₂. Natural silica-based glasses, known as obsidian, are formed when certain volcanic magmas cool rapidly.

Figure 10.32

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Crystalline and amorphous silicon dioxide, SiO₂The difference between crystalline silica and silica glass is shown in these simplified two-dimensional projections. It is readily apparent that much of the ordered arrangement of crystalline silica is lost in the glassy form. But the glass retains enough Si–O bonds to form a hard, rigid material.

Ordinary glass is composed mostly of SiO₂, which usually exists in nature in a crystalline form known as quartz. If quartz (in the form of sand) is melted and allowed to cool, it becomes so viscous that the molecules are unable to move to the low potential energy positions they would occupy in the crystal lattice, so that the disorder present in the liquid gets “frozen into” the solid. In a sense, glass can be regarded as a supercooled liquid. Glasses are transparent because the distances over which disorder appears are small compared to the wavelength of visible light, so there is nothing to scatter the light and produce cloudiness.

Ordinary glass is made by melting silica sand to which has been added some calcium and sodium carbonates. These additives reduce the melting point and make it more difficult for the SiO₂ molecules to arrange themselves into crystalline order as the mixture cools. See here for a brief discussion of this important topic.

[image from Wikimedia Commons]

Figure 10.33

6.  Types of molecular units in solids

Molecules

Molecules, not surprisingly, are the most common building blocks of pure substances. Most of the 15-million-plus chemical substances presently known exist as distinct molecules.

Chemists commonly divide molecular compounds into “small” and “large-molecule” types, the latter usually falling into the class of polymers (see below.) The dividing line between the two categories is not very well defined, and tends to be based more on the properties of the substance and how it is isolated and purified.

Atoms

We usually think of atoms as the building blocks of molecules, so the only pure substances that consist of plain atoms are those of some of the elements — mostly the metallic elements, and also the noble-gas elements. The latter do form liquids and crystalline solids, but only at very low temperatures.

Although the metallic elements form crystalline solids that are essentially atomic in nature, the special properties that give rise to their “metallic” nature puts them into a category of their own.

Most of the non-metallic elements exist under ordinary conditions as small molecules such as O₂ or S₆, or as extended structures that can have a somewhat polymeric nature. Many of these elements can form more than one kind of structure, each one stable under different ranges of temperature and pressure. Multiple structures of the same element are known as allotropes, although the more general term polymorph is now preferred.

Ions

Ions, you will recall, are atoms or molecules that have one or more electrons missing (positive ions) or in excess (negative ions), and therefor possess an electric charge.

A basic law of nature, the electro neutrality principle, states that bulk matter cannot acquire more than a trifling (and chemically insignificant) net electric charge. So one important thing to know about ions is that in ordinary matter, whether in the solid, liquid, or gaseous state, any positive ions must be accompanied by a compensating number of negative ions.

Ionic solids

Ionic solids are covered in much more detail in this section.

Ionic substances such as sodium chloride form crystalline solids that can be regarded as made of ions. These solids tend to be quite hard and have high melting points, reflecting the strong forces between oppositely-charged ions. Solid metal oxides, such as CaO and MgO which are composed of doubly-charged ions don’t melt at all, but simply dissociate into the elements at very high temperatures.