Crystalline solids fall into one of four categories.
| Type of Solid | Intermolecular Forces | Properties | Examples |
|---|---|---|---|
| Ionic | Ionic | High Melting Point, Brittle, Hard | NaCl, MgO |
| Molecular | Hydrogen Bonding, Dipole-Dipole, 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 |
| Covalent Network | Covalent Bonding | High Melting Point, Hard, Nonconducting | C (diamond), SiO2 (quartz) |
The first three categories (ionic, molecular, and metallic solids) all involve packing discrete molecules or atoms into a regular lattice. These three categories are distinguished by the nature of the forces holding the discrete molecules or atoms together. In ionic and molecular solids, there are no chemical bonds between the molecules or atoms. In metallic solids, the metallic bonding between atoms is delocalized over many atoms, producing an electronic band structure.
Network solids, however, contain no discrete molecular units. The atoms in the network solid are held together by conventional covalent bonds with neighboring atoms. The result is a single extended network or array.
Two common examples of network solids are diamond (a form of pure carbon) and quartz (silicon dioxide).
Carbon exists as a pure element at room temperature in three different forms: graphite (the most stable form), diamond, and fullerene.
DiamondThe 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). 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.) Questions 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. |
GraphiteThe 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. Questions 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? |
FullereneUntil 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. 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. Questions 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 graphite? 3. A sample of solid 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? |
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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 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. |
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 sigma bond between the central atom (C or Si) and oxygen. As it happens, the Si-O sigma bond is actually somewhat stronger (452 kJ/mole) than a C-O sigma 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 pi bonds between the carbon and oxygen atoms. These pi bonds provide a home for the electrons on the carbon. By sharing the electrons with the oxygens to form pi 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 2px orbitals and between 2py 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 sigma bonds with four oxygen atoms than it is forming sigma bonds and pi 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 pi bonding orbitals formed from the corresponding py orbitals. Notice that the pi 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, which produces a much more stable structure.
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Carbon Dioxide Pi Bond |
Silicon Dioxide Pi Bond |
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