Atomic arrangement overlapping

We encounter a different type of bond in a nitrogen molecule, N2. There is a single electron in each of the three 2p-orbitals on each atom (33). However, when we try to pair them and form three bonds, only one of the three orbitals on each atom can overlap end to end to form a (T-bond (Fig. 3.10). Two of the 2/7-orbitals on each atom (2px and 2py) are perpendicular to the internuclear axis, and each one contains an unpaired electron (Fig. 3.11, top). When the electrons in one of these p-orbitals on each N atom pair, the orbitals can overlap only in a side-by-side arrangement. This overlap results in a TT-bond, a bond in which the two electrons lie in two lobes, one on each side of the internuclear axis (Fig. 3.11, bottom). More formally, a 7T-bond has a single nodal plane containing the internuclear axis. Although a TT-bond has electron density on each side of the internuclear axis, it is only one bond, with the electron cloud in the form of two lobes, just as a p-orbital is one orbital with two lobes. In a molecule with two Tr-bonds, such as N2, the... 

We have already explained. In terms of hybridisation, how a carbon atom can form four sp hybrid orbitals (see p. 47). We can apply this concept to explain the bonding in alkanes. Ethane is taken as an example of a typical alkane. The four sp hybrid orbitals on each carbon atom will overlap end-on with four other orbitals three hydrogen Is orbitals and one sp hybrid orbital on the other carbon atom. Four cr bonds will be formed and they will adopt a tetrahedral arrangement. This is illustrated for ethane in the diagram. 

Diamond, the hardest of natural materials, consists of a lattice of carbon atoms arranged in a tetrahedral slruclure at equal distances apait (1.544 A) and bonded by electron pairs in localized molecular orbitals formed by overlapping of Ihe. /> hybrids. See aiticle on Diamond. 

Then there would be two Be orbitals available for bonding. This description, however, is still not fully consistent with experimental fact. The Be 2r and 2p orbitals could not overlap a Cl 3p orbital with equal effectiveness that is, this promoted pure atomic arrangement would predict two nonequivalent Be—Cl bonds. Yet we observe experimentally that the Be—Cl bonds are identical in bond length and bond strength. 

Each C atom in C2Hg has four regions of high electron density. The VSEPR theory tells us that each C atom has tetrahedral electronic geometry the resulting atomic arrangement around each C atom has one C and three H atoms at the corners of this tetrahedral arrangement. The VB interpretation is that each C atom is sp hybridized. The C—C bond is formed by overlap of a half-filled sp hybrid orbital of one C atom with a half-filled sp hybrid orbital of the other C atom. Each C—H bond is formed by the overlap of a half-filled sp hybrid orbital on C with the half-filled Ir orbital of an H atom. 

VB theory explains that a covalent bond forms when two atomic orbitals overlap and two electrons with paired (opposite) spins occupy the overlapped region. Orbital hybridization allows us to explain how atomic orbitals mix and change their characteristics during bonding. Based on the observed molecular shape (and the related electron-group arrangement), we postulate the type of hybrid orbital needed. 

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