Triple bonds orbital overlap

As portrayed m Figure 2 20 the two carbons of acetylene are connected to each other by a 2sp-2sp cr bond and each is attached to a hydrogen substituent by a 2sp-ls CT bond The unhybndized 2p orbitals on one carbon overlap with their counterparts on the other to form two rr bonds The carbon-carbon triple bond m acetylene is viewed as a multiple bond of the ct + rr + rr type... 

Section 2 21 Carbon is sp hybridized m acetylene and the triple bond is of the ct + Tt + Tt type The 2s orbital and one of the 2p orbitals combine to give two equivalent sp orbitals that have their axes m a straight line A ct bond between the two carbons is supplemented by two tr bonds formed by overlap of the remaining half filled p orbitals... 

Acetylene is linear and alkynes have a linear geometry of their X—C=C—Y units The carbon-carbon triple bond m alkynes is com posed of a CT and two tt components The triply bonded carbons are sp hybridized The ct component of the triple bond contains two electrons m an orbital generated by the overlap of sp hybndized orbitals on adja cent carbons Each of these carbons also has two 2p orbitals which over lap m parrs so as to give two tt orbitals each of which contains two electrons... 

The degree of overlap of these orbitals is smaller than in the triple bond of an alkyne... 

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. 

A carbon-carbon double bond is stronger than one carbon-carbon single bond but weaker than the sum of two single bonds (Section 2.15). A carbon-carbon triple bond is weaker than the sum of three carbon-carbon single bonds. Recall that a single C—C bond is a o-bond, but the additional bonds in a multiple bond are TT-bonds. One reason for the difference in strength is that the side-by-side overlap of p-orbitals that results in a rr-bond is not as great as the end-to-end overlap that results in a o-bond. 

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. 

FIGURE 1.7 Overlap of p orbitals in a triple bond. For clarity, the a orbitals have been... 

In this chapter, we develop a model of bonding that can be applied to molecules as simple as H2 or as complex as chlorophyll. We begin with a description of bonding based on the idea of overlapping atomic orbitals. We then extend the model to include the molecular shapes described in Chapter 9. Next we apply the model to molecules with double and triple bonds. Then we present variations on the orbital overlap model that encompass electrons distributed across three, four, or more atoms, including the extended systems of molecules such as chlorophyll. Finally, we show how to generalize the model to describe the electronic structures of metals and semiconductors. 

Many of the Lewis structures in Chapter 9 and elsewhere in this book represent molecules that contain double bonds and triple bonds. From simple molecules such as ethylene and acetylene to complex biochemical compounds such as chlorophyll and plastoquinone, multiple bonds are abundant in chemistry. Double bonds and triple bonds can be described by extending the orbital overlap model of bonding. We begin with ethylene, a simple hydrocarbon with the formula C2 H4. 

The triple bond includes two itt bonds formed by the pair-wise overlap of the p orbitals that remain on the carbon and nitrogen atoms. Two figures from different perspectives help illustrate the triple bond ... 

When two p orbitals overlap in a side-by-side configuration, they form a pi bond, shown in Figure 7.7. This bond is named after the Greek letter 7t. The electron clouds in pi bonds overlap less than those in sigma bonds, and they are correspondingly weaker. Pi bonds are often found in molecules with double or triple bonds. One example is ethene, commonly known as ethylene, a simple double-bonded molecule (Figure 7.8). The two vertical p orbitals form a pi bond. The two horizontal orbitals form a sigma bond. 

Double and triple bonds, particularly those in carbon molecules, are often described in MO terms. Thus a double bond is described as consisting of a a bond formed by the overlap of an sp hybrid orbital on each carbon atom and a tt bond formed by the sideways overlap of either the 2p or 2p orbitals (Figure 3.18). A cr orbital has cylindrical symmetry like an atomic s orbital whereas a ir orbital, like an atomic p orbital, has a planar node passing through the nucleus of each of the bonded atoms. A triple bond is similarly described as consisting of a cr orbital and two ir orbitals formed from both the 2p and 2pv orbitals on... 

The left-most C atom (in the structure drawn below) is sp3 hybridized, and the C-H bonds to that C atom are between the sp3 orbitals on C and the Is orbital on H. The other two C atoms are sp hybridized. The right-hand C-H bond is between the sp orbital on C and the Is orbital on H. The c a C triple bond is composed of one sigma bond formed by overlap of sp orbitals, one from each C atom, and two pi bonds, each formed by the overlap of two 2p orbitals, one from each C atom (that is a 2py—2py overlap and a 2pz—2pz overlap). 

In ethylene, there are two types of bonds. Sigma (tr) bonds have the overlap of the orbitals on a line between the two atoms involved in the covalent bond. In ethylene, the C-H bonds and one of the C-C bonds are sigma bonds. Pi (ir) bonds have the overlap of orbitals above and below a line through the two nuclei of the atoms involved in the bond. A double bond is always composed of one sigma and one pi bond. A carbon-to-carbon triple bond results from the... 

For the acetylene molecule, the sp hybrid orbital is formed by combining one 2s orbital and one 2p orbital to form two sp orbitals. Each C atom uses one sp orbital to bond to one Is from H, and the other sp orbital to bond to the other C. The C=C triple bond is formed by the addition of two tt bonds by the overlap of two 2p orbitals from each C. Some of the properties of these hybrid bonds are shown in table 4.13. 

Sigma bonds form when s or p orbitals overlap in a head-on manner. Single bonds cire usually sigma bonds. Pi bonds cire usually double or triple bonds. Figure 5-9 depicts these situations. 

Covalent bonds are formed when atomic orbitals overlap. The overlap of atomic orbitals is called hybridization, and the resulting atomic orbitals are called hybrid orbitals. There are two types of orbital overlap, which form sigma (cr) and pi (tt) bonds. Pi bonds never occur alone without the bonded atoms also being joined by a ct bond. Therefore, a double bond consists of a O bond and a tt bond, whereas a triple bond consists of a ct bond and two tt bonds. A sigma overlap occurs when there is one bonding interaction that results from the overlap of two s orbitals or an s orbital overlaps a p orbital or two p orbitals overlap head to head. A tt overlap occurs only when two bonding interactions result from the sideways overlap of two parallel p... 

---