2.17   Summary

New Concepts

In this chapter, we developed a bonding scheme for alkanes. We continue, as in Chapter 1, to form bonds through the overlap of atomic orbitals. We describe a model in which the atomic orbitals of carbon are combined to form new, hybrid atomic orbitals. The four orbitals resulting from a combination of three 2p orbitals and one 2s orbital are called sp³ hybrid atomic orbitals, reflecting the 75% p character and 25% s character in the hybrid. These hybrid orbitals solve the problems encountered in formation of bonds between pure atomic orbitals. The sp³ hybrids are asymmetric and have a fat and a thin lobe. Overlap between a hydrogen 1s orbital and the fat lobe provides a stronger bond than that between hydrogen 1s and carbon 2p orbitals. In addition, these hybrid orbitals are directed toward the corners of a tetrahedron, which keeps the electrons in the bonds as far apart as possible, thus minimizing destabilizing interactions. Other hybridization schemes are sp², in which the central atom is bonded to three other atoms, and sp, in which the central atom is bonded to two other atoms. Simple molecules in which a central carbon is hybridized sp² are planar, with the three attached groups at the corners of an equilateral triangle. Molecules with sp-hybridized carbons are linear.

Even simple molecules have complicated structures. Methane is a perfect tetrahedron, but we need only substitute a methyl group for one hydrogen of methane for complexity to arise. In ethane, for example, we must consider the consequences of rotation about the carbon–carbon bond. The minimum energy conformation for ethane is the arrangement in which all carbon–hydrogen bonds are staggered. About 3 kcal/mol (12.6 kJ/mol) above this minimum-energy form is the eclipsed form, the transition state, which is the high-energy point (not an energy minimum, but an energy maximum) separating two staggered forms of ethane. The 3 kcal/mol constitutes the rotational barrier between the two staggered forms. This barrier is small compared to the available thermal energy at room temperature, and rotation in ethane is fast.

This chapter again discusses the concept of Lewis acids and bases. The familiar Brønsted bases compete for a proton, but nucelophiles are far more versatile. A nucleophile (Lewis base) is defined as an electron pair donor, and an electrophile (Lewis acid) is anything that reacts with a Lewis base. An electrophile is an electron pair acceptor. We are paid back for these very general definitions with an ability to see as similar all manner of seemingly different reactions. These concepts will run through the entire book.

Key Terms

Reactions, Mechanisms, and Tools

In this chapter, new molecules are built up by first constructing a generic substituted molecule such as methyl―X (CH₃―X). The new hydrocarbon is then generated by letting X = CH₃. In principle, each different carbon–hydrogen bond in a molecule could yield a new hydrocarbon when X = CH₃. In practice, this procedure is not so simple. The problem lies in seeing which hydrogens are really different. The two-dimensional drawings are deceptive. You really must see the molecule in three dimensions before the different carbon–hydrogen bonds can be identified with certainty.

The coding or drawing procedures can get quite abstract. It is vital to be able to keep in mind the real, three-dimensional structures of molecules even as you write them in two-dimensional code. The Newman projection, an enormously useful device for representing molecules three-dimensionally, is described in this chapter.

The naming convention for alkanes is introduced. There are several common or trivial names that are often used and which, therefore, must be learned.

Both ¹H NMR and ¹³C NMR spectroscopy are introduced in this chapter. At this point, the NMR spectrometer functions basically as a machine to determine the numbers of different hydrogens or carbons in a molecule. That function is important, however, both in these early stages of your study of chemistry and in the “real world” of structure determination.

Common Errors

It is easy to become confused by the abstractions used to represent molecules on paper or blackboards. Do not trust the flat surface! It is easy to be fooled by a “carbon” that does not really exist or to be bamboozled by the complexity introduced by ring structures. One good way to minimize such problems (none of us ever really becomes free of the necessity to think about structures presented on the flat surface) is to work lots of isomer problems and play with models.