The Shapes of Some Simple Molecules

So far our concern has emphasized electron bookkeeping. We now turn our attention to the shapes of molecules. 

The tetrahedral geometry of methane is often explained in terms of the valence shell electron-pair repulsion (VSEPR) model. The VSEPR model rests on the idea that an electron pair, either a bonded pair or an unshared pair, associated with a particular atom will be as far away from the atom s other electron pairs as possible. Thus, a tetrahedral geometry permits the four bonds of methane to be maximally separated and is characterized by H—C—H angles of 109.5°, a value referred to as the tetrahedral angle. 

So far we have emphasized stmcture in terms of electron bookkeeping. We now turn our attention to molecular geometry and will see how we can begin to connect the three-dimensional shape of a molecule to its Lewis formula. Table 1.7 lists some simple compounds illustrating the geometries that will be seen most often in our study of organic chemistry. 

Compound Structural formula Repulsive electron pairs Arrangement of electron pairs Molecular shape Molecular model 

Methane (CH4) 109.5 H FI t 109.5° 109.5° 109.5° Carbon has four bonded pairs Tetrahedral Tetrahedral c 

Water (H2O) HvJ Oxygen has two bonded pairs + two unshared pairs Tetrahedral Bent  

Ammonia (NH3) / H Nitrogen has three bonded pairs + one unshared pair Tetrahedral Trigonal pyramidal c 

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When the central atom in a molecule is surrounded only by bonded electron pairs (not by lone pairs) the molecule has a regular geometry or shape which depends on the number of bonded electron pairs. Referring to the molecule as AB, (for convenience sake, the molecule is made up of only two elements A and B) where A is the central atom and x has integral values 2, 3 etc. Table 1.7 gives the arrangement of bonded electron pairs about a central atom and the geometry of some simple molecules. 

Let us now apply these rules and predict the shapes of some molecules. As examples, we are going to use simple molecules we have constructed already in Lecture 2. 

The shapes of the monomeric molecules of the Group 2 halides (gas phase or matrix isolation) pose some interesting problems for those who are content with simple theories of bonding and molecular geometry. Thus, as expected on the basis of either sp hybridization or the... 

Another difficulty with the infrared method is that of determining the band center with sufficient accuracy in the presence of the fine structure or band envelopes due to the overall rotation. Even when high resolution equipment is used so that the separate rotation lines are resolved, it is by no means always a simple problem to identify these lines with certainty so that the band center can be unambiguously determined. The final difficulty is one common to almost all methods and that is the effect of the shape of the potential barrier. The infrared method has the advantage that it is applicable to many molecules for which some of the other methods are not suitable. However, in some of these cases especially, barrier shapes are likely to be more complicated than the simple cosine form usually assumed, and, when this complication occurs, there is a corresponding uncertainty in the height of the potential barrier as determined from the infrared torsional frequencies. In especially favorable cases, it may be possible to observe so-called hot bands i.e., v = 1 to v = 2, 2 to 3, etc. This would add information about the shape of the barrier. 

In this section, we construct a model of molecular shape empirically, which means that we base it on rules suggested by experimental observations rather than on more fundamental principles. We proceed in three steps. First, we set up the basic nodel for simple molecules without lone pairs on the central atom. Then, we elude the effects of lone pairs. Finally, we explore some of the consequences of ecular shape. 

The Lewis stmcture of a molecule shows how its valence electrons are distributed. These stmctures present simple, yet information-filled views of the bonding in chemical species, hi the remaining sections of this chapter, we build on Lewis stmctures to predict the shapes and some of the properties of molecules. In Chapter 10. we use Lewis stmctures as the starting point to develop orbital overlap models of chemical bonding. 

The reader will probably be familiar with at least one simple, qualitative explanation for the shape of the water molecule (there are several). The VSEPR approach (see Section 1.4) predicts that the bond angle should be somewhere between 90° and 109.5°, which is good enough for the purposes of most inorganic chemists. The fact that some of the underlying assumptions in this and other simple theories can be challenged does not necessarily vitiate the theory. 

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