Materials Science

Materials Science Bonding in Solids The Structures of Solids

Materials Science

Have you ever wondered why a silver mirror reflects light that passes through a sheet of glass? Why the metal body of a car left in the summer sun feels so much hotter than the grass next to which it was parked? Why a glass shatters when dropped on the floor, but not the spoon used to stir its contents? Why metals conduct both heat and electricity, but not ceramics such as the tiles that protect the Space Shuttle as it re-enters the atmosphere?

The answers to these questions, and a host of others, lie within the realm of the field known as materials science. The struggle to harness materials isn't new, it traces back to pre-historic times. But recent advances in the synthesis and uses of materials based on an understanding of their structures and properties ­ rather than trial and error ­ have produced new materials with revolutionary properties. We now have ceramic dishes that can go directly from the refrigerator to a hot burner without breaking. Integrated circuits that have brought computers to a point only dreamed of twenty years ago - a computer on every desk. We have new alloys for use in airplanes or tennis rackets that provide strength without weight. And we have taken the first steps toward superconducting materials that will allow us to build trains that float above the surface of the tracks on which they run.

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Bonding in Solids

The first step toward understanding these advances in materials science is to develop a basic understanding of the structure of solids.

Chemists commonly classify solids as either metallic, ionic, molecular, or network covalent on the basis of macroscopic differences in their physical properties, which result from differences in bonding on the atomic scale.

The bonding between atoms in a solid is determined by a combination of two factors: the magnitude of the electronegativities of the atoms in the solid and the differences between these electronegativities. As a result, the bond-type triangle shown in the figure below can be used to predict the classification in which a solid should fall. Compounds with metallic bonds should be metallic solids, those with ionic bonds should be ionic solids, and those with covalent bonds should be either molecular or network covalent solids.

Bond-type triangle showing the classification of bonding types based on average electronegativities and electronegativity differences.

Metallic solids are expected for combinations of atoms whose electronegativities and electronegativity differences place them in the lower left corner of the bond-type triangle shown in the above figure. Ionic compounds are found in the center, toward the top of this triangle, and covalent materials are found in the lower right corner.

Solids are described in terms of their lattice points, which are defined as the regular periodic arrangement of particles in the solid. Metallic solids are composed of planes of individual metal atoms held together by the strong force of attraction between these atoms, as shown in the figure below. Metal atoms don't have enough electrons to fill their valence shells by sharing electrons with their immediate neighbors. Valence electrons are therefore shared by many atoms, delocalizing these electrons over many atoms. Because these valence electrons aren't tightly bound to individual atoms, they are free to migrate through the metal. As a result, metals are good conductors of electricity. Electrons that enter the metal at one edge can displace other electrons to give rise to a net flow of electricity through the metal.

electrons arranged in rows and columns

When the solid contains two or more elements for which there is a large electronegativity difference, electrons are transferred from one atom to another to form an ionic solid. The lattice points in these solids are positive and negative ions, such as the Na+ and Cl- ions shown in the figure below, which are held together by strong ionic bonds.

Na and Cl ions alternating arranged in a cube

Covalent compounds usually crystallize as molecular solids, in which the lattice points are occupied by individual molecules, such as the structure of dry ice shown in the figure below. The individual molecules are held together by relatively strong intramolecular bonds between the atoms that form the molecule. The much weaker intermolecular bonds between these molecules result from the relatively weak van der Waals forces. The van der Waals forces holding CO2 molecules together in dry ice, for example, are so weak that dry ice sublimes --it passes directly from the solid to the gas phase -- at -78ēC.

CO2 molecules with carbon between the O's also arranged in a cube

Network covalent solids are a unique class of materials that can be viewed as a single giant molecule made up of an almost endless number of covalent bonds, such as the structure of diamond shown in the figure below. Because all of the bonds in this structure are equally strong, covalent solids are often very hard and they are notoriously difficult to melt. Diamond is the hardest natural substance and it melts at 3550ēC.

structure of diamond which is in a cube

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The Structure of Solids

The structure of a material is determined by the way in which the simplest repeating units are arranged in space. When these simple repeating units occur in an regular fashion from one end of the solid to the other the material is said to be crystalline. The relationship between the physical properties of a material and its structure is best illustrated by the different elemental forms of carbon.

Carbon occurs as a variety of allotropes. There are two crystalline forms -- diamond and graphite-- and a number of amorphous (noncrystalline) forms, such as charcoal, coke, and carbon black.

References to the characteristic hardness of diamond (from the Greek adamas, "invincible") date back at least 2600 years. It was not until 1797, however, that Smithson Tennant was able to show that diamonds consist solely of carbon. The properties of diamond are remarkable. It has one of the highest melting points (MP = 3550oC) and boiling points (BP = 4827oC) of any naturally occurring substance. It is also the hardest natural substance and it expands less on heating than any other material.

The characteristic properties of diamond are a direct result of its structure. Carbon, with four valence electrons, forms covalent bonds to four neighboring carbon atoms arranged toward the corners of a tetrahedron, as shown in the figure below. Each of these atoms is then bound to four other carbon atoms, which form bonds to four other atoms, and so on. As a result, a perfect diamond is a single giant molecule. The strength of the individual C--C bonds and their arrangement in space give rise to the unusual properties of diamond.

perfect diamond structure

In some ways, the properties of graphite are like those of diamond. Both compounds boil at 4827oC, for example. But graphite is also very different from diamond. Graphite is much less dense than diamond. Whereas diamond is the hardest substance known, graphite is one of the softest. Diamond is an excellent insulator, graphite is such a good conductor of electricity that graphite electrodes are used in electrical cells.

The physical properties of graphite can be understood from the structure of the solid shown in the figure below. Graphite consists of extended planes of carbon atoms in which each carbon forms strong covalent bonds to three other carbon atoms. (The strong bonds between carbon atoms within each plane explain the exceptionally high melting point and boiling point of graphite.) These planes of atoms, however, are held together by relatively weak van der Waals forces. Because the bonds between planes of atoms are weak, it is easy to deform the solid by allowing one plane of atoms to move relative to another. Graphite is therefore soft enough to be used in pencils and as a lubricant in motor oil.

graphite arranged in planes

The characteristic properties of graphite and diamond might lead you to expect that diamond would be more stable than graphite. This isn't what is observed experimentally. Graphite at 25C and 1 atm pressure is slightly more stable than diamond. (The enthalpy of atom combination of graphite is -5557 kJ/molrxn and that of diamond is -5523 kJ/molrxn). At very high temperatures and pressures, however, diamond becomes more stable than graphite. In 1955 General Electric developed a process to make industrial-grade diamonds by treating graphite with a metal catalyst at temperatures of 2000 to 3000 K and pressures above 125,000 atm. Roughly 40% of industrial-quality diamonds are now synthetic. Although gem-quality diamonds can be synthesized, the costs involved are prohibitive.

Both diamond and graphite occur as regularly packed crystals. Other forms of carbon are amorphous --they lack a regular structure. Charcoal results from heating wood in the absence of oxygen. To make carbon black, natural gas is burned in a limited amount of air to give a thick, black smoke that contains extremely small particles of carbon that can be collected when the gas is cooled and passed through an electrostatic precipitator. Coke is a more regularly structured material, closer in structure to graphite than either charcoal or carbon black, which is made from coal.

In 1985 a third stable, crystalline form of carbon was made by vaporizing graphite with a laser. The product of this reaction is a molecule with the formula C60 that has a structure with the symmetry of a soccer-ball. Because this structure resembles the geodesic dome invented by R. Buckminster Fuller, C60 was named buckminsterfullerene, or "buckyball" for short.

Some of the fascination of C60 can be understood by contrasting this form of pure carbon with diamond and graphite. C60 is unique because it exists as distinct molecules, not extended arrays of atoms. Equally important, C60 can be obtained as a pure substance, whereas the surfaces of diamond and graphite are inevitably contaminated by hydrogen atoms that bind to the carbon atoms on the surface.

C60 is now known to be a member of a family of compounds known as the fullerenes. C60 may be the most important of the fullerenes because it is the most perfectly symmetric molecule possible, spinning in the solid at a rate of more than 100 million times per second. Because of their symmetry, C60 molecules pack as regularly as ping-pong balls. The resulting solid has unusual properties. Initially, it is as soft as graphite, but when compressed by 30%, it becomes harder than diamond. When this pressure is released, the solid springs back to its original volume. C60 therefore has the remarkable property that it bounces back when shot at a metal surface at high speeds.

C60 also has the remarkable ability to form compounds in which it is an insulator, a conductor, a semiconductor, or a superconductor. By itself, C60 is a semiconductor. When mixed with just enough potassium to give a compound with the empirical formula K3C60, it conducts electricity like a metal. When excess potassium is added, this solid becomes an insulator. When K3C60 is cooled to 18 K, the result is a superconductor. The potential of fullerene chemistry for both practical materials and laboratory curiosities is large enough to explain why this molecule has been described as "exocharmic" --it exudes charm.

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