Tetrahedral covalent bonding

The selenium-selenium and tellurium-tellurium distances observed in the manganese compounds provide further evidence for this structural interpretation. They correspond to radii that agree more satisfactorily with the normal-valence radii than with the tetrahedral radii of the nonmetallic atoms, indicating that the atoms are not forming tetrahedral covalent bonds ... 

This mode of transition from one type of bond to the other allows the resistance of a purely ionic bond to be taken as 1.0 (bond by s orbitals), and the resistance of a tetrahedral covalent bond (sp3 and roughly sp and sp2), as 2.0 (Table 3.6). Now, adding to unity the degree of covalence (Table 3.5) expressed in tenths and hundredths, we obtain the magnitudes of specific bond resistance for an arbitrary state staying in the range 1.0-2.0. The specific bond resistance thus calculated, e.g., between A1 and O is for A1203 1.00 + 0.41 = 1.41, and for Zn and S, 1.00 +0.76 = 1.76. 

The structure of dislocations in diamond lattices, which, as discussed in Chap. 3, is adopted quite frequently by elements that have tetrahedral covalent bonding, has to conform to the comparatively rigid tetrahedral... 

The bond ionicity in phosphides and diphosphides can be estimated as follows. We suggested earlier a tetrahedral covalent bond scheme for compounds [8]. This scheme is... 

An electron-valence scheme of bonds in CrSi2 is proposed, based on the assumption that the tetravalent chromium atoms are bonded to the four nearest silicon atoms, lying in adjacent hexagonal layers, by somewhat distorted d s tetrahedral covalent bonds, and the silicon atoms are bonded by sp hybrid bonds to two chromium atoms and two silicon atoms. 

Hybridization of the available Ss and 5p orbitals of tin produces a suitable situation for the formation of four tetrahedral covalent bonds Many tin compounds such as the tm(TV) halides (except fluoride) and the organotins have this covalent mode of bonding. 

Unlike the forces between ions which are electrostatic and without direction, covalent bonds are directed in space. For a simple molecule or covalently bonded ion made up of typical elements the shape is nearly always decided by the number of bonding electron pairs and the number of lone pairs (pairs of electrons not involved in bonding) around the central metal atom, which arrange themselves so as to be as far apart as possible because of electrostatic repulsion between the electron pairs. Table 2.8 shows the essential shape assumed by simple molecules or ions with one central atom X. Carbon is able to form a great many covalently bonded compounds in which there are chains of carbon atoms linked by single covalent bonds. In each case where the carbon atoms are joined to four other atoms the essential orientation around each carbon atom is tetrahedral. 

The two kinds of covalent bond are not identical, one being a simple covalent bond, a sigma (ct) bond, the other being a stronger (but more reactive) bond called a n bond (p. 56). As in the formation of methane both elements attain noble gas configurations. We can consider the formation of ethene as the linking of two tetrahedral carbon atoms to form the molecule C2H4 represented as ... 

The formation of a fourth covalent bond by the aluminium atom results in spatial rearrangement from the trigonal planar, for three bonding electron pairs, to tetrahedral, for four bonding electron pairs. 

The unequal distribution of charge produced when elements of different electronegativities combine causes a polarity of the covalent bond joining them and, unless this polarity is balanced by an equal and opposite polarity, the molecule will be a dipole and have a dipole moment (for example, a hydrogen halide). Carbon tetrachloride is one of a relatively few examples in which a strong polarity does not result in a molecular dipole. It has a tetrahedral configuration... 

All the other aluminium halides are covalently bonded with aluminium showing a coordination number of four towards these larger halogen atoms. The four halogen atoms arrange themselves approximately tetrahedrally around the aluminium and dimeric molecules are produced with the configuration given below ... 

The H—O—H angle m water (105°) and the H—N—H angles m ammonia (107°) are slightly smaller than the tetrahedral angle These bond angle contractions are easily accommodated by VSEPR by reasoning that electron pairs m bonds take up less space than an unshared pair The electron pair m a covalent bond feels the attractive force of... 

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

In all the groups along the chain, the bond angle is fixed. It is determined by considering a carbon atom at the centre of a regular tetrahedron and the four covalent bonds are in the directions of the four comers of the tetrahedron. This sets the bond angle at 109° 28 as shown in Fig. A.4 and this is called the tetrahedral angle. 

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). 

We have constructed a number of sets of atomic radii for use in compounds containing covalent bonds. These radii have been obtained from the study of observed interatomic distances. They are not necessarily applicable only to crystals containing pure covalent bonds (it is indeed probable that very few crystals of this type exist) but also to crystals and molecules in which the bonds approach the covalent type more closely than the ionic or metallic type. The crystals considered to belong to this class are tetrahedral crystals, pyrite and marcasite-type crystals, and others which have been found on application of the various criteria discussed in the preceding section to contain covalent bonds or bonds which approach this extreme. 

In other crystals an octahedral metal atom is attached to six non-metal atoms, each of which forms one, two, or three, rather than four, bonds with other atoms. The interatomic distance in such a crystal should be equal to the sum of the octahedral radius of the metal atom and the normal-valence radius (Table VI) of the non-metal atom. This is found to be true for many crystals with the potassium chlorostannate (H 61) and cadmium iodide (C 6) structures (Table XIB). Data are included in Table XIC for crystals in which a tetrahedral atom is bonded to a non-metal atom with two or three covalent bonds. The values of dcalc are obtained by adding the tetrahedral radius for the former to the normal-valence radius for the latter atom. 

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