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Network Solids

Crystalline Solids

Crystalline solids fall into one of four categories.

Type of Solid Interaction Properties Examples
Ionic Ionic High Melting Point, Brittle, Hard NaCl, MgO
Molecular Hydrogen Bonding,
London Dispersion
Low Melting Point, Nonconducting H2, CO2
Metallic Metallic Bonding Variable Hardness and Melting Point (depending upon strength of metallic bonding), Conducting Fe, Mg
Network Covalent Bonding High Melting Point, Hard, Nonconducting C (diamond),
SiO2 (quartz)

All four categories involve packing discrete molecules or atoms into a lattice or repeating array, though network solids are a special case. The categories are distinguished by the nature of the interactions holding the discrete molecules or atoms together.

In ionic and molecular solids, there are no chemical bonds between the molecules, atoms, or ions. The solid consists of discrete chemical species held together by intermolecular forces that are electrostatic or coulombic in nature. This behavior is most obvious for an ionic solid such as NaCl, where the positively charged Na+ ions are attracted to the negatively charged Cl- ions. Even in the absence of ions, however, electrostatic forces are operational. For polar molecules such as CH2Cl2, the positively charged region of one molecular is attracted to the negatively charged region of another molecule (dipole-dipole interactions). For a nonpolar molecule such as CO2, which has no permanent dipole moment, the random motion of electrons gives rise to temporary polarity (a temporary dipole moment). Electrostatic attractions between two temporarily polarized molecules are called London Dispersion Forces.

Hydrogen bonding is a term describing an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation. (See the IUPAC Provisional Recommendation on the definition of a hydrogen bond.) Dots are employed to indicate the presence of a hydrogen bond: X–H···Y. The attractive interaction in a hydrogen bond typically has a strong electrostatic contribution, but dispersion forces and weak covalent bonding are also present.

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. (A note about bonding.)

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.

Network Solids

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. Two common examples of network solids are diamond (a form of pure carbon) and quartz (silicon dioxide). In quartz one cannot detect discrete SiO2 molecules. Instead the solid is an extended three-dimensional network of ...-Si-O-Si-O-... bonding.


Carbon exists as a pure element at room temperature in three different forms: graphite (the most stable form), diamond, and fullerene.


The structure of diamond is shown at the right in a "ball-and-stick" format. The balls represent the carbon atoms and the sticks represent a covalent bond. Be aware that in the "ball-and-stick" representation the size of the balls do not accurately represent the size of carbon atoms. In addition, a single stick is drawn to represent a covalent bond irrespective of whether the bond is a single, double, or triple bond or requires resonance structures to represent. In the diamond structure, all bonds are single covalent bonds (sigma bonds). The "space-filling" format is an alternate representation that displays atoms as spheres with a radius equal to the van der Waals radius, thus providing a better sense of the size of the atoms. Bonds are not drawn in the space-filling depiction.

Notice that diamond is a network solid. The entire solid is an "endless" repetition of carbon atoms bonded to each other by covalent bonds. (In the display at the right, the structure is truncated to fit in the display area.)


  1. What is the bonding geometry around each carbon?
  2. What is the hybridization of carbon in diamond?
  3. The diamond structure consists of a repeating series of rings. How many carbon atoms are in a ring?
  4. Diamond are renowned for its hardness. Explain why this property is expected on the basis of the structure of diamond.


The most stable form of carbon is graphite. Graphite consists of sheets of carbon atoms covalently bonded together. These sheets are then stacked to form graphite. The display at the right shows a ball-and-stick representation of graphite. The sheets extended "indefinitely" in the xy plane, but the structure has been truncated for display purposed.

Graphite may also be regarded as a network solid, even though there is no bonding in the z direction. Each layer, however, is an "endless" bonded network of carbon atoms.


  1. What is the bonding geometry around each carbon?
  2. What is the hybridization of carbon in graphite?
  3. The a layer of the graphite structure consists of a repeating series of rings. How many carbon atoms are in a ring?
  4. What force holds the carbon sheets together in graphite?
  5. Graphite is very slippery and is often used in lubricants. Explain why this property is expected on the basis of the structure of graphite.
  6. The slipperiness of graphite is enhanced by the introduction of impurities. Where would such impurities be located and why would they make graphite a better lubricant?


Until the mid 1980's, chemistry textbooks stated that pure carbon existed in two forms: graphite and diamond. The discovery of C60 molecules in interstellar dust in 1985 added a third form to this list. The existence of C60, which resembles a soccer ball, had been hypothesized by theoretians for many years. In the late 1980's synthetic methods were developed for the synthesis of C60, and the ready availability of this form of carbon led to extensive research into its properties.

The C60 molecule, shown at the right in ball-and-stick form, is called buckminsterfullerene, though the shorter name fullerene is often used. The name is a tribute to the American architect R. Buckminster Fuller, who is famous for designing and constructing geodesic domes which bear a close similarity to the structure of C60. As is evident from the display, C60 is a sphere composed of six-member and five-member carbon rings. These balls are sometimes fondly referred to as "Bucky balls".

It should be noted that fullerenes are an entire class of pure carbon compounds rather than a single compound. A distorted sphere containing more than 60 carbon atoms have also been found, and it is also possible to create long tubes. All of these substances are pure carbon.

It should also be noted that fullerences are molecular solids rather than network solids. The various fullerenes exist as individual molecules, not an extended structure with continuous bonding.


  1. What is the bonding geometry around each carbon? (Note that this geometry is distorted in C60.)
  2. What is the hybridization of carbon in fullerene?
  3. A single crystal of C60 falls into which class of crystalline solids?
  4. It has been hypothesized that C60 would make a good lubricant. Why might C60 make a good lubricant?

Silicon Dioxide

Silicon dioxide (SiO2), also called silica, occurs naturally in many forms. Quartz is essentially pure silicon dioxide. Sand is composed of small quartz fragments. Many precious gems are quartz containing colored impurities. Amethyst is quartz colored red by the presence of iron(III) ions. Agate and onyx are also quartz containing impurities. Flint is silica colored black by carbon.

Quartz has a very complicated crystal structure, which involves interwoven helical chains. When heated to about 1500o C, quartz changes into the mineral cristobalite, whose structure is shown at the right in ball-and-stick form. The brown balls represent the silicon atoms and the red balls represent the oxygen atoms. Cristobalite is pure SiO2. Notice the similarity in structure between cristobalite and diamond.

All silicates involve silicon in tetrahedral environments surrounded by oxygen atoms. Observe how Cristobalite is an example of a network solid. There are no discrete SiO2 units in this compound. Chemical bonding between SiO2 units leads to an extended network silicon-oxygen bonding.

CO2 and SiO2

Carbon and silicon are both Group IV nonmetals and thus have very similar electronic structures. Under ambient conditions, however, carbon dioxide is a gas and silicon dioxide is a hard network solid.

Why are the physical properties of CO2 and SiO2 so different?

Both compounds involve a σ bond between the central atom (C or Si) and oxygen. As it happens, the Si-O σ bond is actually somewhat stronger (452 kJ/mole) than a C-O σ bond (358 kJ/mole). In the CO2 molecule, however, there is excellent overlap between the carbon 2px and 2py orbitals and the oxygen 2px and 2py orbitals, with the result being the formation of very strong π bonds between the carbon and oxygen atoms. These π bonds provide a home for the electrons on the carbon. By sharing the electrons with the oxygens to form π bonds, no other bonding interactions are possible for the carbon. Thus CO2 exists as discrete molecules, which just happen to be nonpolar owing to the geometry of the molecule. Consequently the only intermolecular forces in pure carbon dioxide are London dispersion forces.

The electron structure of SiO2 is very different. Unlike CO2, where there is excellent overlap between carbon and oxygen 2p orbitals, there is very poor overlap between the silicon 3px and 3py orbitals and the oxygen 2px and 2py orbitals. The net result is the silicon is much more stable forming σ bonds with four oxygen atoms than it is forming σ bonds and π bonds with two oxygen atoms. This behavior produces a network solid rather than discrete SiO2 molecules.

The electron density plots of the yz plane, shown below, show π bonding orbitals formed from the corresponding py orbitals. Notice that the π bond in CO2 involves excellent orbital overlap with substantial electron density in the region in between the carbon and oxygen nuclei. In contrast, there is very poor overlap in the SiO2 molecule. The electron density plot on the right was calculated for an isolated SiO2 molecule, where the silicon atom was only able to interact with two oxygen atoms. In silica, each silicon atom has the opportunity to interact with four oxygen atoms (σ bonding), which produces a much more stable structure.

Carbon Dioxide π Bond
Silicon Dioxide π Bond

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