Hybrid Orbitals and the Structure of Ethylene

Problem 1.9 Convert the following molecular model of hexane a component of gasoline, into a 

Although sp hybiidization is the most common electronic state of carbon, it s not the only possibility. Look at ethylene, C2H4, for example. It was recognized more than 100 years ago that ethylene carbons can be tetravalent only if they share four electrons and are linked by a double bond. Furthermore, ethylene is planar (flat) and has bond angles of approximately 120° rather than 109.5°. 

Predicting the Structures of Simple Organic Molecules from Their Formulas 

Commonly used in biology as a tissue preservative, formaldehyde, CH2O, contains a carbon-r x 5c// double bond. Draw the line-bond structure of formaldehyde, and indicate the hybridization of the carbon atom. 

To complete the structure of ethylene, four hydrogen atoms form a bonds with the remaining four sp orbitals. Ethylene thus has a planar structure, with H-C-H and H-C-C bond angles of approximately 120°. (The actual values are 117.4° for the H-C-H bond angle and 121.3° for the H-C-C bond angle.) Each C-H bond has a length of 108.7 pm and a strength of 465 kj/mol (111 kcal/mol). 

Draw a line-bond structure for propane, CH3CH2CH3. Predict the value of each bond angle, and indicate the overall shape of the molecule. 

The bonds we ve seen in methane and ethane are called single bonds because they result from the sharing of one electron pair between bonded atoms. It was recognized nearly 150 years ago, however, that carbon atoms can also form double bonds by sharing two electron pairs between atoms or triple bonds by sharing three electron pairs. Ethylene, for instance, has the structure H2C = CH2 and contains a carbon-carbon double bond, while acetylene has the stmcture HC CH and contains a carbon-carbon triple bond. 

How are multiple bonds described by valence bond theory When we discussed sp hybrid orbitals in Section 1.6, we said that the four valence-shell atomic orbitals of carbon combine to form four equivalent sp hybrids. Imagine instead that the 2s orbital combines with only two of the three available 

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Atomic Structure The Nucleus Atomic Structure Orbitals 4 Atomic Structure Electron Configurations 6 Development of Chemical Bonding Theory 7 The Nature of Chemical Bonds Valence Bond Theory sp Hybrid Orbitals and the Structure of Methane 12 sp Hybrid Orbitals and the Structure of Ethane 13 sp2 Hybrid Orbitals and the Structure of Ethylene 14 sp Hybrid Orbitals and the Structure of Acetylene 17 Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur 18 The Nature of Chemical Bonds Molecular Orbital Theory 20 Drawing Chemical Structures 21 Summary 24... 

Rule 3 If two or three pairs of electrons form a multiple bond between two atoms, the first bond is a sigma bond formed by a hybrid orbital. The second bond is a pi bond, consisting of two lobes above and below the sigma bond, formed by two unhybridized p orbitals (see the structure of ethylene in Figure 2-17). The third bond of a triple bond is another pi bond, perpendicular to the first pi bond (shown in Figure 2-18). 

The structure of ethylene and the orbital hybridization model for its double bond were presented m Section 2 20 and are briefly reviewed m Figure 5 1 Ethylene is planar each carbon is sp hybridized and the double bond is considered to have a a component and a TT component The ct component arises from overlap of sp hybrid orbitals along a line connecting the two carbons the tt component via a side by side overlap of two p orbitals Regions of high electron density attributed to the tt electrons appear above and below the plane of the molecule and are clearly evident m the electrostatic potential map Most of the reactions of ethylene and other alkenes involve these electrons... 

Figure 1.14 The structure of ethylene. Orbital overlap of two sp hybridized carbons forms a carbon-carbon double bond. One part of the double bond results from a (head-on) overlap of sp2 orbitals (green), and the other part results from (sideways) overlap of unhybridized p orbitals (red/blue). The ir bond has regions of electron density on either side of a line drawn between nuclei.

The structure of ethylene and the orbital hybridization model for the double bond were presented in Section 1.17. To review. Figure 5.1 depicts the planar structure of ethylene, its bond distances, and its bond angles. Each of the carbon atoms is xp -hybridized, and the double bond possesses a o component and a tt component. The o component results when an sp orbital of one carbon, oriented so that its axis lies along the intemuclear axis, overlaps with a similarly disposed sp orbital of the other carbon. Each sp orbital contains one electron, and the resulting a bond contains two of the four electrons of the double bond. The tt bond contributes the other two electrons and is formed by a side-by-side overlap of singly occupied p orbitals of the two carbons. 

The structure of ethylene, C2H4, is shown in Fig. 8-2. The molecule is planar, and each carbon is bonded to two hydrogens and to the other carbon, With three groups attached to each carbon, we use a set of s-f hybrid orbitals for v bonding. 

Let s now examine compounds with a double bond. Wherever there is a double bond, sp hybridization should be considered for the atoms involved. For example, second-period elements use a combination of an sp hybrid orbital and the unhybridized 2p atomic orbital to form double bonds. Consider ethylene, C2H4, whose Lewis structure is shown in Figure 1.20(a). A cr bond between the carbons... 

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)). 

The great advantage of this method is that it can be used to build up structures of much larger molecules quickly and without having to imagine that the molecule is made up from isolated atoms. So it is easy to work out the structure of ethene (ethylene) the simplest alkene. Ethene is a planar molecule with bond angles dose to 120°. Our approach will be to hybridize all the orbitals needed for the C-H framework and see what is left over. In this case we need three bonds from each carbon atom (one to make a C-C bond and two to make C-H bonds). 

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