Diamond is one of the hardest-known substances, whereas graphite is soft enough to be used as pencil lead. These very different properties stem from the different arrangements of the carbon atoms in the different allotropes. You may be less familiar with a recently discovered form of carbon: graphene. Graphene was first isolated in by using tape to peel off thinner and thinner layers from graphite. It is essentially a single sheet one atom thick of graphite.
Graphene, illustrated in Figure 8 , is not only strong and lightweight, but it is also an excellent conductor of electricity and heat.
These properties may prove very useful in a wide range of applications, such as vastly improved computer chips and circuits, better batteries and solar cells, and stronger and lighter structural materials. In a crystalline solid, the atoms, ions, or molecules are arranged in a definite repeating pattern, but occasional defects may occur in the pattern. Several types of defects are known, as illustrated in Figure 9. Vacancies are defects that occur when positions that should contain atoms or ions are vacant.
Less commonly, some atoms or ions in a crystal may occupy positions, called interstitial sites , located between the regular positions for atoms. Other distortions are found in impure crystals, as, for example, when the cations, anions, or molecules of the impurity are too large to fit into the regular positions without distorting the structure.
Trace amounts of impurities are sometimes added to a crystal a process known as doping in order to create defects in the structure that yield desirable changes in its properties. For example, silicon crystals are doped with varying amounts of different elements to yield suitable electrical properties for their use in the manufacture of semiconductors and computer chips.
Some substances form crystalline solids consisting of particles in a very organized structure; others form amorphous noncrystalline solids with an internal structure that is not ordered. The main types of crystalline solids are ionic solids, metallic solids, covalent network solids, and molecular solids.
The properties of the different kinds of crystalline solids are due to the types of particles of which they consist, the arrangements of the particles, and the strengths of the attractions between them. Because their particles experience identical attractions, crystalline solids have distinct melting temperatures; the particles in amorphous solids experience a range of interactions, so they soften gradually and melt over a range of temperatures.
Some crystalline solids have defects in the definite repeating pattern of their particles. These defects which include vacancies, atoms or ions not in the regular positions, and impurities change physical properties such as electrical conductivity, which is exploited in the silicon crystals used to manufacture computer chips.
Ice has a crystalline structure stabilized by hydrogen bonding. These intermolecular forces are of comparable strength and thus require the same amount of energy to overcome.
The force of attraction between atoms in metals, such as copper and aluminum, or alloys, such as brass and bronze, are metallic bonds. Molecular, ionic, and covalent solids all have one thing in common. With only rare exceptions, the electrons in these solids are localized.
They either reside on one of the atoms or ions or they are shared by a pair of atoms or a small group of atoms. Metal atoms don't have enough electrons to fill their valence shells by sharing electrons with their immediate neighbors. Electrons in the valence shell are therefore shared by many atoms, instead of just two. In effect, the valence electrons are delocalized over many metal atoms. Because these electrons aren't tightly bound to individual atoms, they are free to migrate through the metal.
What makes diamond so hard? Let's compare it to a few of the softer minerals in the scale. Gypsum is a crystalline hydrate of calcium sulfate: CaSO 4. Calcite is a mineral form of calcium carbonate: CaCO 3. Fluorite is composed of calcium fluoride: CaF 2. These materials are all ionic solids. Gypsum contains calcium cations and sulfate anions, as well as bound water.
Calcite contains calcium ions and carbonate ions. Fluorite contains calcium ions and fluoride ions. Surely the ionic bonds that hold these ions together are very strong.
Why can they be deformed and scratched? An ionic bond is not directional. It does not matter whether a sulfate ion is above or below a calcium ion, or to the left or the right of it. The attraction between the ions is still the same. If we imagine a shearing force on a crystal of calcium sulfate, meaning that we are pushing against only one layer of the crystalline material, then we might see that layer slide in response to the force.
As the ions in that layer slide, they begin to lose their attraction to the ions in the layer beneath them, but then they become attracted to new ions that are sliding towards. When we are finished, we still have an ionic solid in which ions are atracted to counterions around them, but the partners have changed. In metallic solids and network solids, however, chemical bonds hold the individual chemical subunits together.
The crystal is essential a single, macroscopic molecule with continuous chemical bonding throughout the entire structure. In metallic solids, the valence electrons are no longer exclusively associated with a single atom. Instead these electrons exist in molecular orbitals that are delocalized over many atoms, producing an electronic band structure.
The metallic crystal essentially consists of a set of metal cations in a sea of electrons. This type of chemical bonding is called metallic bonding. You learned previously that an ionic solid consists of positively and negatively charged ions held together by electrostatic forces. The strength of the attractive forces depends on the charge and size of the ions that compose the lattice and determines many of the physical properties of the crystal. The lattice energy i.
In both cases, however, the values are large; that is, simple ionic compounds have high melting points and are relatively hard and brittle solids. Molecular solids consist of atoms or molecules held to each other by dipole—dipole interactions, London dispersion forces, or hydrogen bonds, or any combination of these. The arrangement of the molecules in solid benzene is as follows:. For similar substances, the strength of the London dispersion forces increases smoothly with increasing molecular mass.
For example, the melting points of benzene C 6 H 6 , naphthalene C 10 H 8 , and anthracene C 14 H 10 , with one, two, and three fused aromatic rings, are 5. The enthalpies of fusion also increase smoothly within the series: benzene 9. If the molecules have shapes that cannot pack together efficiently in the crystal, however, then the melting points and the enthalpies of fusion tend to be unexpectedly low because the molecules are unable to arrange themselves to optimize intermolecular interactions.
Self-healing rubber is an example of a molecular solid with the potential for significant commercial applications. The material can stretch, but when snapped into pieces it can bond back together again through reestablishment of its hydrogen-bonding network without showing any sign of weakness. Among other applications, it is being studied for its use in adhesives and bicycle tires that will self-heal. Covalent solids are formed by networks or chains of atoms or molecules held together by covalent bonds.
A perfect single crystal of a covalent solid is therefore a single giant molecule. The carbon atoms form six-membered rings. The unit cell of diamond can be described as an fcc array of carbon atoms with four additional carbon atoms inserted into four of the tetrahedral holes.
It thus has the zinc blende structure described in Section Elemental silicon has the same structure, as does silicon carbide SiC , which has alternating C and Si atoms. The structure of crystalline quartz SiO 2 , shown in Section All compounds with the diamond and related structures are hard, high-melting-point solids that are not easily deformed.
Instead, they tend to shatter when subjected to large stresses, and they usually do not conduct electricity very well. It is difficult to deform or melt these and related compounds because strong covalent C—C or Si—Si or polar covalent Si—C or Si—O bonds must be broken, which requires a large input of energy. Other covalent solids have very different structures. It contains planar networks of six-membered rings of sp2 hybridized carbon atoms in which each carbon is bonded to three others.
To completely describe the bonding in graphite, we need a molecular orbital approach similar to the one used for benzene in Chapter 9. In fact, the C—C distance in graphite In graphite, the two-dimensional planes of carbon atoms are stacked to form a three-dimensional solid; only London dispersion forces hold the layers together.
As a result, graphite exhibits properties typical of both covalent and molecular solids. It is also very soft; the layers can easily slide past one another because of the weak interlayer interactions. Finally, graphite is black because it contains an immense number of alternating double bonds, which results in a very small energy difference between the individual molecular orbitals. Thus light of virtually all wavelengths is absorbed. Diamond, on the other hand, is colorless when pure because it has no delocalized electrons.
In network solids, conventional chemical bonds hold the chemical subunits together. The bonding between chemical subunits, however, is identical to that within the subunits, resulting in a continuous network of chemical bonds.
One common examples of network solids are diamond a form of pure carbon Carbon exists as a pure element at room temperature in three different forms: graphite the most stable form , diamond, and fullerene.
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