Triple bonds molecular shape

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. 

Number of Bonded Pairs on the Central Atom (Double and Triple Bonds Count Only as 1 Pair) Number of Lone Electron Pairs on the Central Atom Molecular Shape (Molecular Geometry)... 

The characterization of the interrelations between chemical bonding and molecular shape requires a detailed analysis of the electronic density of molecules. Chemical bonding is a quantum mechanical phenomenon, and the shorthand notations of formal single, double, triple, and aromatic bonds used by chemists are a useful but rather severe oversimplification of reality. Similarly, the classical concepts of body and surface , the usual tools for the shape characterization of macroscopic objects, can be applied to molecules only indirectly. The quantum mechanical uncertainty of both electronic and nuclear positions within a molecule implies that valid descriptions of both chemical bonding and molecular shape must be based on the fuzzy, delocalize properties of electronic density distributions. These electron distributions are dominated by the nuclear arrangements and hence quantum mechanical uncertainly affects electrons on two levels by the lesser positional uncertainty of the more massive nuclei, and by the more prominent positional uncertainty of the electrons themselves. These two factors play important roles in chemistry and affect both chemical bonding and molecular shape. 

The bonding pi orbital 71) follows regions separate from a line drawn between the two atoms in a bond. Two overlapping p orbitals will form n bonds to contain the additional shared electrons in molecules with double or triple bonds, n bonds prevent atoms from rotating about the central axis between them. The atomic orbitals that form the compound C2H4 are shown above to the left. Each H atom contains one electron in an s orbital and each C atom contains 4 valence electrons in three hybrid sp2 orbitals and one p orbital. The compound itself contains the molecular orbitals shown below to the left. There are five a bonds (white ovals) and one n bond (the shaded shapes). For additional examples of n bonds in carbon compounds, see Skill 6.1a. 

Carbon is able to form sp, sp2, and sp3 hybridized atomic orbitals and a and n molecular orbitals (see Skill 1.3c). For example, in CH4, the electron density of the four sp3 orbitals of C each overlap with an s orbital of H to form four a bonds. In C2H4 (an alkene), two sp2 orbitals on each C overlap with H, the remaining sp2 orbitals overlap with each other in a crbond, and the p orbitals (drawn as shaded shapes) overlap with each other above and beneath the carbons in a n bond (also drawn as shaded shapes). In CO2, the C atom has two sp hybrid orbitals and two p orbitals. These form one crbond and one n bond with the two unfilled p orbitals on each 0 atom. In C2H2 (an alkyne), a triple bond forms with one crand two n. 

Molecular packing in the low-temperature orthorhombic phase of perdeutero-acetylene shows a layer structure with a T-shaped interaction between pairs of molecules within a layer. The D—C distances form an isosceles triangle with edges of 2.738, 2.737 A and a D to C=C midpoint distance of 2.672 A. The D—C contacts are some 0.2 A shorter than those expected from van der Waals radii considerations. The geometrical arrangement clearly implies a donor-acceptor interaction between D and the triple-bond TT-electrons, a situation described as long ago as 1972 to account for close C=C—I interactions in crystals of diiodo-acetylene. More recently five O—H—C C and five NH—C=C TT-interactions with mean H—C distances of 2.69 (for O—H) and 2.61 (for N—H) have been located in the CSD. 

Ethene, H2C = CH2, serves as the starting material for the synthesis of polyethylene, from which plastic bags and milk jugs are made. Ethyne, H—C=C—H, is used as a fuel for welding torches. The double bond in ethene and the triple bond in ethyne have the same effect on molecular shape as single bonds. Predict the shapes and bond angles of ethene and ethyne. 

The problem with Lewis dot structures is that they provide no insight into molecular shapes, orbitals, or distributions of electrons within molecules. Instead, they are only useful for predicting the number of bonds an atom forms whether the atom has lone pairs and whether single, double, or triple bonds are used. Once an atom is found to have an octet using a Lewis analysis, no further insight into the structure or reactivity can be obtained from the Lewis structure. We have to turn to more sophisticated molecular structure and bonding concepts to understand structure and reactivity. 

Molecular Shapes The shapes of molecules can be predicted by combining Lewis theory with valence shell electron pair repulsion (VSEPR) theory. In tiiis model, electron groups— lone pairs, single bonds, double bonds, and triple bonds—aroxmd the central atom repel one another and determine the geometry of the molecule. 

The properties of molecules are directly related to their shapes. In VSEPR theory, molecular geometries are determined by the repulsions between electron groups on the central atom. An electron group can be a single bond, double bond, triple bond, lone pair, or even a single electron. 

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