Acetylene hybrid atomic orbitals

Similar, but different, redeployment is envisaged when a carbon atom combines with three other atoms, e.g. in ethene (ethylene) (p. 8) three sp2 hybrid atomic orbitals disposed at 120° to each other in the same plane (plane trigonal hybridisation) are then employed. Finally, when carbon combines with two other atoms, e.g. in ethyne (acetylene) (p. 9) two sp1 hybrid atomic orbitals disposed at 180° to each other (idigonal hybridisation) are employed. In each case the s orbital is always involved as it is the one of lowest energy level. 

A similar description applies to a linear ethyne (acetylene) molecule, H-OC-H. Now the carbon atoms are sp hybridized, and the o bonds are built from hybrid atomic orbitals of the form... 

Many of the reactions in which acetylene participates, as well as many properties of acetylene, can be understood in terms of the stmcture and bonding of acetylene. Acetylene is a linear molecule in which two of the atomic orbitals on the carbon are sp hybridized and two are involved in 7T bonds. The lengths and energies of the C—H O bonds and C=C<7 + 27t bonds are as follows ... 

Comparison of the electronic structures of acetonitrile and propyne (methylacetylene). In both compounds, the atoms at the ends of the triple bonds are sp hybridized, and the bond angles are 180°. In place of the acetylenic hydrogen atom, the nitrile has a lone pair of electrons in the sp orbital of nitrogen,... 

Figure 8-8 The acetylene molecule, C2H2. (a) The overlap diagram of two rp-hybridized carbon atoms and two r orbitals from two hydrogen atoms. The hybridized sp orbitals on each C are shown in green and the unhybridized p orbitals are shown in tan. The dashed lines, each connecting two lobes, indicate the side-by-side overlap of the four unhybridized p orbitals to form two tr bonds. There are two C—H cr bonds, one C—C cr bond (green, hatched), and two C—C it bonds (hatched and cross-hatched). This makes the net carbon-carbon bond a triple bond, (b) The tr bonding orbitals tan) are positioned with one above and below the Une of the cr bonds (green) and the other behind and in front of the line of the cr bonds.

Because of the greater proportion of s character in the sp o orbitals of the carbon atoms than that in sp hybrids (alkenes) or sp hybrids (alkanes), any bonding pair of electrons can come closer to carbon nuclei in acetylene than in alkenes or alkanes. Hence hydrogen is released as a proton more easily in the acetylenic carbon atoms. The increased acidity of hydrogen attached to a carbon atom with a triple bond results in the formation of a class of carbides with certain metals, called acetylides. These acetylides are unstable and highly sensitive to shock and heat, exploding violently (see Chapter 30). 

Figure 1.22 shows a Lewis structure and an orbital overlap diagram for acetylene, C2H2. A carbon-carbon triple bond consists of one sigma bond and two pi bonds. The sigma bond is formed by the overlap of sp hybrid orbitals. One pi bond is formed by the overlap of a pair of parallel 2p atomic orbitals. The second pi bond is formed by the overlap of a second pair of parallel 2p atomic orbitals. 

The carbon atom of an acetylene is connected to only two other atoms. Therefore, we combine the 2s orbital with only one 2p orbital to make two sp-hybiid orbitals (Figure 3. 15). These orbitals extend in opposite directions from the carbon atom. The angle between the two hybrid orbitals is 180° so as to minimize repulsion between any electrons placed in them. One valence electron is placed in each sp-hybrid orbital. The remaining two valence electrons occupy two different p orbitals that are perpendicular to each other and perpendicular to the hybrid sp orbitals. 

Because each carbon m acetylene is bonded to two other atoms the orbital hybridization model requires each carbon to have two equivalent orbitals available for CT bonds as outlined m Figure 2 19 According to this model the carbon 2s orbital and one of Its 2p orbitals combine to generate two sp hybrid orbitals each of which has 50% s character and 50% p character These two sp orbitals share a common axis but their major lobes are oriented at an angle of 180° to each other Two of the original 2p orbitals remain unhybridized... 

In the third type of hybridisation of the valence electrons of carbon, two linear 2sp orbitals are formed leaving two unhybridised 2p orbitals. Linear a bonds are formed by overlap of the sp hybrid orbitals with orbitals of neighbouring atoms, as in the molecule ethyne (acetylene) C2H2, Fig. 1, A3. The unhybridised p orbitals of the carbon atoms overlap to form two n bonds the bonds formed between two C atoms in this way are represented as Csp Csp, or simply as C C. 

When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap head-on to form a strong sp-sp a bond. In addition, the pz orbitals from each carbon form a pz-pz it bond by sideways overlap and the py orbitals overlap similarly to form a py-py tt bond. The net effect is the sharing of six electrons and formation of a carbon-carbon triple bond. The two remaining sp hybrid orbitals each form a bond with hydrogen to complete the acetylene molecule (Figure 1.16). 

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. 

In triple-bond compounds (e.g., acetylene), carbon is connected to only two other atoms, and hence uses sp hybridization, which means that the four atoms are in a straight line (Fig. 1.6). Each carbon has two p orbitals remaining, with one electron in each. These orbitals are perpendicular to each other and to the C—C axis. They overlap in the manner shown in Figure 1.7 to form two n orbitals. A triple bond is thus composed of one a and two n orbitals. Triple bonds between carbon and nitrogen can be represented in a similar manner. 

Each carbon atom has a steric number of 2, indicating that acetylene is a linear molecule and that sp hybrid orbitals can be used to construct the bonding orbital framework. Figure 10-24a shows the a bonding system of acetylene. 

Both carbon atoms in the acetylene molecule undergo sp hybridization. Two p orbitals remain unhybridized. So, one sp hybrid orbital from each carbon atom overlaps with the s orbital of a hydrogen atom and two C — Ho bonds result. Also, between the two adjacent C atoms a C—Co bond is formed as a result of end to end overlap of the sp hybrid orbitals. So in the C2H2 molecule there are three o bonds in total. 

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