Hybrid orbitals structures

Fig. 2.2. Several representations of ethylene, which has two bonded carbon atoms, (a) is a ball and stick model, (b) shows a chemical structure, and (c) is a schematic orbital diagram. Note the trigonal planar arrangement of the bonds between the atoms-this arrangement is the consequence of the hybrid orbital structure.

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... 

As proven in Chapter 13.Ill, this two-configuration description of Be s electronic structure is equivalent to a description is which two electrons reside in the Is orbital (with opposite, a and (3 spins) while the other pair reside in 2s-2p hybrid orbitals (more correctly, polarized orbitals) in a manner that instantaneously correlates their motions ... 

The structure of ethylene and the orbital hybridization model for its double bond were presented m Section 2 20 and are briefly reviewed m Figure 5 1 Ethylene is planar each carbon is sp hybridized and the double bond is considered to have a a component and a TT component The ct component arises from overlap of sp hybrid orbitals along a line connecting the two carbons the tt component via a side by side overlap of two p orbitals Regions of high electron density attributed to the tt electrons appear above and below the plane of the molecule and are clearly evident m the electrostatic potential map Most of the reactions of ethylene and other alkenes involve these electrons... 

In most metals the electron behaves as a particle having approximately the same mass as the electron in free space. In the Group IV semiconductors, dris is usually not the case, and the effective mass of electrons can be substantially different from that of the electron in free space. The electronic sUmcture of Si and Ge utilizes hybrid orbitals for all of the valence elecU ons and all electron spins are paired within this structure. Electrons may be drermally separated from the elecU on population in dris bond structure, which is given the name the valence band, and become conduction elecU ons, creating at dre same time... 

For c/rwo-boranes and for the larger open-cluster boranes it becomes increasingly difficult to write a simple satisfactory localized orbital structure, and a full MO treatment is required. Intermediate cases, such a.s B5H9, require several resonance hybrids in the localized orbital... 

Humulene. structure of, 202 Hund s rule, 6 sp Hybrid orbitals. 17-18 sp2 Hybrid orbitals, 15. sp3 Hybrid orbitals, 12-14 Hydrate, 701... 

Strategy The first step is to draw Lewis structures. To find the hybridization of nitrogen, place in hybrid orbitals—... 

Radicals with very polar substituents e.g. trifluoromethyl radical 2), and radicals that arc part of strained ring systems (e.g. cydopropyl radical 3) arc ct-radicals. They have a pyramidal structure and are depicted with the free spin resident in an spJ hybrid orbital. nr-Radicals with appropriate substitution are potentially chiral, however, barriers to inversion are typically low with respect to the activation energy for reaction. 

Self-Test 3.7B Suggest a structure in terms of hybrid orbitals for each carbon atom in ethyne, C2H2. 

Now consider the alkynes, hydrocarbons with carbon-carbon triple bonds. The Lewis structure of the linear molecule ethyne (acetylene) is H—O C- H. To describe the bonding in a linear molecule, we need a hybridization scheme that produces two equivalent orbitals at 180° from each other this is sp hybridization. Each C atom has one electron in each of its two sp hybrid orbitals and one electron in each of its two perpendicular unhybridized 2p-orbitals (43). The electrons in the sp hybrid orbitals on the two carbon atoms pair and form a carbon—carbon tr-bond. The electrons in the remaining sp hybrid orbitals pair with hydrogen Ls-elec-trons to form two carbon—hydrogen o-bonds. The electrons in the two perpendicular sets of 2/z-orbitals pair with a side-by-side overlap, forming two ir-honds at 90° to each other. As in the N2 molecule, the electron density in the o-bonds forms a cylinder about the C—C bond axis. The resulting bonding pattern is shown in Fig. 3.23. 

Account for the structure of a formic acid molecule, HCOOH, in terms of hybrid orbitals, bond angles, and cr- and TT-bonds. The C atom is attached to an H atom, a terminal O atom, and an —OH group. 

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