Chemistry covalent crystals

From the early days of structural chemistry there has been considerable interest in discussing bond lengths in terms of radii assigned to the elements, and it has become customary to do this in terms of three sets of radii, applicable to metallic, ionic, and covalent crystals. Distances between non-bonded atoms have been compared with sums of van der Waals radii , assumed to be close to ionic radii. The earliest covalent radii for non-metals were taken as one-half of the M—M distances in molecules or crystals in which M forms % — N bonds N being the number of the Periodic Group), that is, from molecules such as F2, HO-OH, H2N--NH2, P4, Sg, and the crystalline elements of Group IV with the diamond structure. This accounts... 

Transition metal clusters, however, need still to be tested in the engineering of crystalline materials. Crystal engineering has been defined as the capacity to make crystals with a purpose. In transition metal cluster chemistry this purpose is that of utilizing the distinct characteristics mentioned above to construct crystals that can function as the result of the inter-cluster interactions. To do this the experimentalist needs to conceive ways of directing the crystal-building process towards given architectures, i. e. needs to learn how to make non-covalent crystal synthesis. Clearly, the growth and success of a solid-state chemistry of transition metal clusters depends crucially on a close interaction between synthesis, theory, solid state characterization, and evaluation of properties. 

When applied to crystalline solids, the paradigm shift leads directly from supramolecular chemistry to crystal engineering. What is a (molecular) crystal if not an organized entity of higher complexity held together by in-termolecular forces (J.M. Lehn) Who can deny that the collective properties of such a giant supermolecule are the result of the convolution of intermolecular non-covalent bonding between molecular/ionic components with the periodicity of the crystal. 

In order to be able to calculate the concentrations of point defects at thermodynamic equilibrium, it is necessary to know the change in free energy of the crystal which accompanies the formation of point defects, since the equilibrium is determined by the minimization of the free energy when the pressure, the temperature, and the other independent thermodynamic variables are given. A theoretical calculation of the free energy of formation of defects is still one of the most difficult problems in solid state physics and chemistry. The methods of calculation for each group of materials - metals, covalent crystals, ionic crystals - are all very... 

The idea of point defects in crystals goes back to Frenkel, who in 1926 proposed the existence of point defects to explain the observed values of ionic conductivity in crystalline solids. In a crystal of composition MX such as a monovalent metal halide or a divalent metal oxide or sulfide, volume ionic conductivity occurs by motion of positive or negative ions in the lattice under the influence of an electric field. If the crystal were perfect, imperfections, such as vacant lattice sites or interstitial atoms, would need to be created for ionic conductivity to occur. A great deal of energy is required to dislodge an ion from its normal lattice position and thus the current in perfect crystals would be very, very small under normal voltages. To get around this difficulty, Frenkel proposed that point defects existed in the lattice prior to the application of the electric field. This, of course, has been substantiated by subsequent work and the concept of point defects in all classes of solids, metals, ionic crystals, covalent crystals, semiconductors, etc., is an important part of the physics and chemistry of crystalline solids, not only with respect to ionic conductivity but also with respect to diffusion, radiation damage, creep, and many other properties. 

A systematic approach to the crystal chemistry of borides is possible on the simple basis of atom size considerations, as well as the tendency of B to form covalent skeletons. 

Sulfur exists in two crystalline forms, rhombic and monoclinic, the latter comprising three axes of unequal length, two of which intersect at right angles. The bonding within each crystal lattice is covalent and with an electronic structure approaching the configuration of an inert gas atom, the element shows purely nonmetallic chemistry. 

Examples of molecular crystals are found throughout organic, organometallic, and inorganic chemistry. Low melting and boiling temperatures characterize the crystals. We will look at just two examples, carbon dioxide and water (ice), both familiar, small, covalently bonded molecules. 

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