Diamond covalent bonding

J/moFK, whereas the S° of diamond is 2.44 J/moFK. In diamond, covalent bonds extend in three dimensions, allowing the atoms little movement in graphite, covalent bonds extend only within a two-dimensional sheet, and motion of the sheets relative to each other is relatively easy. 

Pure carbon occurs naturally in two modifications, diamond and graphite. In both these forms the carbon atoms are linked by covalent bonds to give giant molecules (Figure S.2). 

Covalent bonding appears in its pure form in diamond, silicon and germanium - all materials with large moduli (that of diamond is the highest known). It is the dominant... 

If the polymer is completely cross-linked (/= 1) then the modulus (Ej) is known it is that of diamond, 10 GPa. If it has no covalent bonds at all, then the modulus (E2) is that of a simple hydrocarbon like paraffin wax, and that, too, is known it is 1 GPa. 

In diamond, each carbon atom forms single bonds with four other carbon atoms arranged tetrahedrally around it The hybridization in diamond is sp3. The three-dimensional covalent bonding contributes to diamond s unusual hardness. Diamond is one of the hardest substances known it is used in cutting tools and quality grindstones (Figure 9.12). 

Even though silicon is metallic in appearance, it is not generally classified as a metal. The electrical conductivity of silicon is so much less than that of ordinary metals it is called a semiconductor. Silicon is an example of a network solid (see Figure 20-1)—it has the same atomic arrangement that occurs in diamond. Each silicon atom is surrounded by, and covalently bonded to, four other silicon atoms. Thus, the silicon crystal can be regarded as one giant molecule. 

FIGURE 5.21 The structure of diamond, Each sphere represents the location of the center of a carbon atom. Each atom is at the center of a tetrahedron formed hy the sp1 hybrid covalent bonds to each of its four neighbors. 

Questions such as, for example, whether sphalerite contains Zn++ and S= ions or has a covalent structure similar to that of diamond, and whether ionic or covalent bonds are present in complexes such as [FeF% —, [Fe(CN)e]=, etc., have been extensively discussed it has, indeed, until recently not been at all clear whether or not they could be definitely... 

Ionic bonds may be fully as strong as covalent bonds, so that properties such as hardness, solubility, melting point, ionization in solution, and chemical character are not especially valuable criteria as a rule. Sometimes comparison of properties with those of compounds of known bond type permits reasonably certain conclusions to be drawn. Thus the similarity in physical properties as well as in atomic arrangement of SiC, AIN, and diamond suggests that all three substances contain covalent bonds. PbS is like FeS2, MoS2, etc. in properties rather than like CaS, so that it is improbable that PbS is an ionic substance. 

Replacement of each water molecule by an MnAli2 icosahedron with shared faces leads to an infinite framework with 136 Mn and 816 Al atoms in the unit cube. This framework is similar to the framework of covalently bonded carbon atoms in a diamond crystal, with one body diagonal of each pentagonal dodecahedron in place of each C—C covalent bond. 

Diamondoids, when in the solid state, melt at much higher temperatures than other hydrocarbon molecules with the same number of carbon atoms in their structures. Since they also possess low strain energy, they are more stable and stiff, resembling diamond in a broad sense. They contain dense, three-dimensional networks of covalent bonds, formed chiefly from first and second row atoms with a valence of three or more. Many of the diamondoids possess structures rich in tetrahedrally coordinated carbon. They are materials with superior strength-to-weight ratio. 

Carbon in the form of diamond is an electrical insulator because of its huge band gap. hi fact, its band gap of 580 kJ/mol substantially exceeds the C—C bond energy of 345 kJ/mol. In other words, it requires more energy to promote an electron from band to band in diamond than to break a covalent bond. Lead, in contrast, is a metallic conductor because it has... 

Network solids such as diamond, graphite, or silica cannot dissolve without breaking covalent chemical bonds. Because intermolecular forces of attraction are always much weaker than covalent bonds, solvent-solute interactions are never strong enough to offset the energy cost of breaking bonds. Covalent solids are insoluble in all solvents. Although they may react with specific liquids or vapors, covalent solids will not dissolve in solvents. 

When covalent bonds favor neighbors of the same element, the positions c and d can also be occupied by atoms of the same kind as in a or b. This applies to diamond and to the Zintl phase NaTl NaTl can be regarded as a network of Tl particles that form a diamond structure which encloses Na+ ions (cf. Fig. 13.3, p. 134). 

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