Hybridization and Bonding in Ethylene

The oxidation number of carbon decreases from —2 in CH3CI to —4 in CH3Li. 

The generalization illustrated by the preceding examples can be expressed in terms broad enough to cover these reactions and many others, as follows Oxidation of carbon occurs when a bond between a carbon and an atom that is less electronegative than carbon is replaced by a bond to an atom that is more electronegative than carbon. The reverse process is reduction. 

Both of the following reactions will be encountered in Chapter 4. One is oxidation-reduction, the other is not. Which is which  

Which of the following reactions requires an oxidizing agent, a reducing agent, or neither  

Sample Solution The CH3 carbon is unchanged in the reaction however, the carbon of CH2OH now has two bonds to oxygen. Therefore, the reaction requires an oxidizing agent. 

Describe the bonding in propane according to the orbital hybridization model. 

In the next few sections we ll examine the application of the valence bond-orbital hybridization model to alkenes and alkynes, then return to other aspects of alkanes in Section 2.11. We ll begin with ethylene. 

Each carbon atom still has, at this point, an unhybridized 2p orbital available for bonding. These two half-filled 2p orbitals have their axes perpendicular to the framework of a bonds of the molecule and overlap in a side-by-side maimer to give a pi (it) bond. The carbon-carbon double bond of ethylene is viewed as a combination of a ct bond plus a it bond. The additional increment of bonding makes a carbon-carbon double bond both stronger and shorter than a carbon-carbon single bond. 

Electrons in a tt bond are called it electrons. The probability of finding a ir electron is highest in the region above and below the plane of the molecule. The plane of the molecule corresponds to a nodal plane, where the probability of finding a tt electron is zero. 

In general, you can expect that carbon will be sp -hybridized when it is directly bonded to three atoms in a neutral molecule. 

FIGURE 1.25 (a) Electron configuration of carbon in its most stable state. (6) An electron is promoted from the 2s orbital to the vacant 2p orbital, (c) The 2s orbital and two of the three 2p orbitals are combined to give a set of three equal-energy sp -hybridized orbitals. One of the 2p orbitals remains unchanged. 

One measure of the strength of a bond is its bond dissociation energy. This topic wiii be introduced in Section 4.17 and appiied to ethyiene in Section 5.2. 

FIGURE 1.26 Representation of orbital mixing in sp hybridization. Mixing of one s orbital with two p orbitals generates three sp hybrid orbitals. Each sp hybrid orbital has one-third s character and two-thirds p character. The axes of the three sp hybrid orbitals are coplanar. One 2p orbital remains unhybridized, and its axis is perpendicular to the plane defined by the axes of the sp orbitals. 

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Methane and the Biosphere 59 Bonding in Ethane 61 sp Hybridization and Bonding in Ethylene 61... 

It turns out, in fact, that the electron distribution and bonding in ethylene can be equally well described by assuming no hybridization at all. The "bent bond" model depicted at the right requires only that the directions of some of the atomic-p orbitals be distorted sufficiently to provide the overlap needed for bonding. So one could well argue that hybrid orbitals are not real they do turn out to be convenient for understanding the bonding of simple molecules at the elementary level, and this is why we use them. 

FIGURE 2 17 The carbon-carbon double bond in ethylene has a cr component and a tt compo nent The cr component arises from overlap of sp hybridized orbitals along the internuclear axis The tt component results from a side by side overlap of 2p orbitals... 

FIGURE 5 1 (a) The planar framework of u bonds in ethylene showing bond distances and angles (b) and (c) The p orbitals of two sp hybridized carbons overlap to produce a tt bond (d) The electrostatic potential map shows a region of high negative potential due to the tt elec trons above and below the plane of the atoms... 

The hybridization concept can also be applied to molecules containing double and triple bonds. The descriptive valence bond approach to the bonding in ethylene and... 

We conclude this introduction to hydrocarbons by describing the orbital hybridization model of bonding in ethylene and acetylene, parents of the alkene and alkyne families, respectively. 

The double bond in ethylene contains one a bond and one 7r bond. The a bond forms from the end-on overlap of two hybrid orbitals, and the 7i bond forms from the side-by-side overlap of two atomic p orbitals. Figure 10-21 shows the complete orbital picture of the bonding in ethylene. Ethylene is the simplest of a class of molecules, the alkenes, all of which contain CDC double bonds. The alkenes are the subject of our Box on page 404. 

FIGURE 16.10 The structure and bonding in the anion of Zeise s salt. A cr-bond results from the overlap of a dsp2 hybrid orbital on the metal and the 7T orbital on ethylene. Back donation from a d orbital on the metal to the -k orbital on ethylene gives some 7r bonding (shown in (c)). 

Figure 10-4 shows the hybridization that occurs in ethylene, H2C=CH2. Each carbon has sp2 hybridization. On each carbon, two of the hybrid orbitals overlap with an s-orbital on a hydrogen atom to form a carbon-to-hydrogen covalent bond. The third sp2 hybrid orbital overlaps with the sp2 hybrid on the other carbon to form a carbon-to-carbon covalent bond. Note that each carbon has a remaining p-orbital that has not undergone hybridization. These are also overlapping above and below a line joining the carbons. 

The carbon-oxygen double bond in aldehydes and ketones is similar and can be described in either of these two ways. If we adopt the iocalised-orbital description, formaldehyde will have two directed lone pairs in place of two of the C-H bonds in ethylene. In this case the axes of these hybrid orbitals will be in the molecular plane (unlike the oxygen lone pairs in water). Either the components of the double bond or the lone pairs can be transformed back into symmetry forms. The alternative description of the lone pairs would he one er-type along the 0-0 direction and one jr-type with axis perpendicular to the 0-0 bond hut in the molecular plane. It is the latter orbital which has the highest energy, so that an electron is removed from it in. ionisation or excitation to the lowest excited state. 

The double bond between carbon and oxygen looks just like the double bond in ethylene. There is a sigma bond formed by overlap of sp2 hybrid orbitals and a pi bond formed by overlap of the unhybridizedp orbitals on carbon and oxygen (Figure 2-19). 

In a molecule such as acetylene with a triple carbon-carbon bond, the first C—C bond and the C—H bond lie along the continuation of each other (sp hybridization), so that the molecule is linear. The second and third bonds are n—n bonds similar to the second bond in ethylene. The bond length at 1.20 A is only 78 % of that of the single bond. 

Atomic orbitals that do not participate in the hybridization are then used for bonding with other atomic orbitals on adjacent centers as long as there is nonzero overlap of the atontic orbitals. For example, a hybridization description for the bonding in ethylene accounts for a so-called sigma (a) framework of bonding utihzing sp hybrids and a second interaction called a pi (tt) interaction between pure p atontic orbitals on the carbon atoms (8). 

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