Oxygen hybrid orbitals

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... 

The oxygen m furan has two unshared electron pairs (Figure 11 16c) One pair is like the pair m pyrrole occupying a p orbital and contributing two electrons to complete the SIX TT electron requirement for aromatic stabilization The other electron pair m furan IS an extra pair not needed to satisfy the 4n + 2 rule for aromaticity and occupies an sp hybridized orbital like the unshared pair m pyridine The bonding m thiophene is similar to that of furan... 

Phosphorus and sulfur are the third-row analogs of nitrogen and oxygen, and the bonding in both can be described using hybrid orbitals. Because of their positions in the third row, however, both phosphorus and sulfur can expand their outer-shell octets and form more than the typical number of covalent bonds. Phosphorus, for instance, often forms five covalent bonds, and sulfur occasionally forms four. 

Both inner atoms have steric numbers of 4 and tetrahedral electron group geometry, so both can be described using s p hybrid orbitals. All four hydrogen atoms occupy outer positions, and these form bonds to the inner atoms through 1 s-s p overlap. The oxygen atom has two lone pairs, one in each of the two hybrid orbitals not used to form O—H bonds. 

Which orbitals hold these lone pairs We answer this question by taking an inventory of the valence orbitals. For the inner atom, the s hybrid orbitals account for the 2 S orbital and the two 2 p orbitals that lie in the plane of the molecule, and the third p orbital, perpendicular to the molecular plane, is part of n system. All of the valence orbitals of the inner oxygen atom are accounted for. For each outer atom, one p orbital in the plane of the page contributes to the a bonds, and the p orbital perpendicular to the plane of the page is part of the n system. The remaining valence orbitals on each outer atom are the 2. S orbital and the 2 orbital that lies in the molecular... 

The molecule has 16 valence electrons. Its Lewis stmcture shows that the molecule has two double bonds, with a steric number of 2 for the carbon atom. Consistent with this, the molecule is linear. Figure 10-39 shows the two a bonds formed by end-on overlap between sp hybrid orbitals on the carbon atom and 2 Pz atomic orbitals of oxygen. 

Figure 3.14 Hybrid orbitals h, and Ii2 formed from the s and p orbitals of the oxygen atom to correspond to the bond angle of 104.5°. These orbitals have a greater overlap with the hydrogen Is orbitals than the atomic 2p orbitals and so form stronger bonds.

An oxygen atom can also form a double bond to carbon thus in propanone (acetone), Me2C=Q , the oxygen atom could use three sp2 hybrid orbitals one to form a a bond by overlap with an sp2 orbital of the carbon atom, and the other two to accommodate the two lone pairs of electrons. This leaves an unhybridised p orbital on both oxygen and carbon, and these can overlap with each other laterally (cf. C=C, p. 9) to form a n bond ... 

The double bond that is shown in each of the two structures just shown is not localized as is reflected by the two resonance structures. However, the two single bonds and the unshared pair are localized as a result of the hybrid orbitals in which they reside. The hybrid orbital type is sp2, which accounts for the bond angle being 119.5°. There is one p orbital not used in the hybridization that is perpendicular to the plane of the molecule, which allows for the tv bonding to the two oxygen atoms simultaneously. The n bond is described as being delocalized, and this can be shown as follows ... 

In dimethyl ether, the oxygen atom is sp3 hybridized. In creating two single bonds, each bond is formed by the overlap of one of its sp3 hybrid orbitals with the sp3 hybrid orbital on the adjacent carbon atom. Each of the remaining two hybrid orbitals on the oxygen atom contain a lone pair of electrons. The resulting molecule is polar. The intermolecular forces found operating between molecules of dimethyl ether are therefore dipole-dipole interactions and London forces. 

Let s begin assigning valence electrons by half-filling three sp2 hybrid orbitals on the sulfur atom (atom A) and one sp2 hybrid orbital on each of the oxygen atoms (atoms B, C and D)... 

Next, we half-fill the lone unhybridized 3p orbital on sulfur and the lone 2p orbital on the oxygen atom with a formal charge of zero (atom B). Following this, the 2p orbital of the other two oxygen atoms (atoms C and D), are filled and then lone pairs are placed in the sp2 hybrid orbitals that are still empty. At this stage, then, all 24 valence electrons have been put into atomic and hybrid orbitals on the four atoms. Now we overlap the six half-filled sp2 hybrid orbitals to generate the cr-bond framework and combine the three 2p orbitals (2 filled, one half-filled) and the 3p orbital (half-filled) to form the four 7t-molecular orbitals, as shown below ... 

It can be imagined that the bonds can be broken at any location, that is, with an oxygen, hydroxy, silicon, or aluminum exposed. In this case, it could further be imagined that s-, p-, and sp3-hybridized orbitals would be on the surface. This... 

Clays contain aluminum oxides in addition to silicon as Si02 and its polymeric forms. Again, there are the p orbitals of oxygen and sp-hybridized orbitals from aluminum, which may result in end-on-end or side-by-side bonding with the same restrictions encountered with silicon. 

Oxygen forms two bonds by utilising its half filled sp3 hybrid orbitals when it forms the H20 molecule with hydrogen. 

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