Lone pair structures

Antimony forms polymeric oxyhalides, and not metallic as in BiOCl. The fluoride, SbOF, has been prepared in two forms V, with a ladder structure, and iM which has a layered structure. Both forms have a trigonal bipyramidal structure about antimony with three oxygens, one fluorine and one lone pair. Structural parameters are given in Table 15, from which it can be seen that L-SbOF heads the table as the nearest to an ideal fit for trigonal bipyramidal geometry. 

The Te 4d level in the compound has shifted by only 1.0 eV to a higher energy (less bound) with respect to its position in the elemental form. The differences in the valence bands of the compound and the element are clearly observed. The sharp peak at 2eV in Te is attributed to the upper nonbonding lone electron pairs associated with twofold coordinated tellurium. In the mixed crystal this peak is absent, suggesting that the lone pair structure typical of the element is lost. The... 

If the spatial arrangement of atoms is required this can be deduced from the basic structure by neglecting the positions occupied by lone pairs of electrons. Water, for example, can be described as a V shape whilst ammonia is a trigonal pyramid. 

The element before carbon in Period 2, boron, has one electron less than carbon, and forms many covalent compounds of type BX3 where X is a monovalent atom or group. In these, the boron uses three sp hybrid orbitals to form three trigonal planar bonds, like carbon in ethene, but the unhybridised 2p orbital is vacant, i.e. it contains no electrons. In the nitrogen atom (one more electron than carbon) one orbital must contain two electrons—the lone pair hence sp hybridisation will give four tetrahedral orbitals, one containing this lone pair. Oxygen similarly hybridised will have two orbitals occupied by lone pairs, and fluorine, three. Hence the hydrides of the elements from carbon to fluorine have the structures... 

This structure indicates that carbon monoxide should have donor properties, the carbon atom having a lone pair of electrons. Carbon... 

Ammonia is a colourless gas at room temperature and atmospheric pressure with a characteristic pungent smell. It is easily liquefied either by cooling (b.p. 240 K) or under a pressure of 8-9 atmospheres at ordinary temperature. Some of its physical and many of its chemical properties are best understood in terms of its structure. Like the other group head elements, nitrogen has no d orbitals available for bond formation and it is limited to a maximum of four single bonds. Ammonia has a basic tetrahedral arrangement with a lone pair occupying one position ... 

Phosphine is a colourless gas at room temperature, boiling point 183K. with an unpleasant odour it is extremely poisonous. Like ammonia, phosphine has an essentially tetrahedral structure with one position occupied by a lone pair of electrons. Phosphorus, however, is a larger atom than nitrogen and the lone pair of electrons on the phosphorus are much less concentrated in space. Thus phosphine has a very much smaller dipole moment than ammonia. Hence phosphine is not associated (like ammonia) in the liquid state (see data in Table 9.2) and it is only sparingly soluble in water. 

The structure of sulphur tetrafluoride, and probably also SeF and TeF4, is trigonal bipyramidal with one position occupied by a lone pair of electrons ... 

In xenon difluoride, the electronic structure shows three lone pairs around the xenon, and two covalent bonds to the two fluorine atoms hence it is believed that here xenon is using one p (doublepear) orbital to form two bonds ... 

Figure 2-51. a) The rotational barrier in amides can only be explained by VB representation using two resonance structures, b) RAMSES accounts for the (albeit partial) conjugation between the carbonyl double bond and the lone pair on the nitrogen atom. 

RAMSES is usually generated from molecular structures in a VB representation. The details of the connection table (localized charges, lone pairs, and bond orders) are kept within the model and are accessible for further processes. Bond orders are stored with the n-systems, while the number of free electrons is stored with the atoms. Upon modification oF a molecule (e.g., in systems dealing with reactions), the VB representation has to be generated in an adapted Form from the RAMSES notation. 

Another example of reduced symmetry is provided by the changes that occur as H2O fragments into OH and H. The a bonding orbitals (ai and b2) and in-plane lone pair (ai) and the a antibonding (ai and b2) of H2O become a orbitals (see the Figure below) the out-of-plane bi lone pair orbital becomes a" (in Appendix IV of Electronic Spectra and Electronic Structure of Polyatomic Molecules, G. Herzberg, Van Nostrand Reinhold Co., New York, N.Y. (1966) tables are given which allow one to determine how particular... 

Fig. 1-6). The structure obtained for thiazoie is surprisingly close to an average of the structures of thiophene (169) and 1,3,4-thiadiazole (170) (Fig. 1-7). From a comparison of the molecular structures of thiazoie, thiophene, thiadiazole. and pyridine (171), it appears that around C(4) the bond angles of thiazoie C(4)-H with both adjacent C(4)-N and C(4)-C(5) bonds show a difference of 5.4° that, compared to a difference in C(2)-H of pyridine of 4.2°, is interpreted by L. Nygaard (159) as resulting from an attraction of H(4) by the electron lone pair of nitrogen. 

The corresponding resonance description shows the delocalization of the nitrogen lone pair electrons m terms of contributions from dipolar structures... 

The reason for the slow hydrolysis compared to that of structurally similar compounds like nitrones or 0,lV-acetals might be the following (b-67MI50800) in the protonated species (77) assistance of the lone pair of electrons at nitrogen is sterically hindered due to the large angle of its orbital to the plane of the ring. 

In acyclic structures, such effects are averaged by rotation, but in cyclic structures differences in C—H bond strengths based on the different alignments can be recognized. The C—H bonds that are in an anti orientation to the lone pair are weaker than the C—H bonds in other orientations. 

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