Of course, the real lesson is the topic of discussion for this chapter: bonding and chemical interactions. We will not actually address complex chemical bonding, such as that which takes place in the Maillard reaction, in this chapter. (We cannot stress enough, however, that the nucleophilic mechanism by which many reactions, like the Maillard reaction, proceed will be tested in the MCAT’s Biological Sciences section.) Rather, this chapter will address the basics of chemical bonding and interactions. Here, we will investigate the nature and behavior of covalent and ionic bonds. We will also review a system by which bonding electrons are accounted for, Lewis structures, and go over the main principles of valence shell electron pair repulsion (VSEPR) theory. Finally, we will recount the various modes of interaction between molecules, the intermolecular forces.
Bonding
The atoms of most elements, except for a few noble gases, can combine to form molecules. The atoms in most molecules are held together by strong attractive forces called chemical bonds, which are formed via the interaction of the valence electrons of the combining atoms. The chemical and physical properties of the resulting compound are usually very different from those of the constituent elements. For example, elemental sodium, an alkali metal, is so reactive that it can actually produce fire when reacting with water (the reaction is highly exothermic), and elemental chlorine gas is so toxic that it was used for chemical warfare during World War I. However, when an atom of sodium and an atom of chlorine react, the produced ionic compound, sodium chloride, is safe for us to eat. You may know it better as common table salt!
Bridge
Electronegativity (which we learned about in the last chapter) is a property that addresses how an individual atom acts within a bond and will help us understand the quality of the molecules formed from atoms with different electronegativities.
Why do atoms join together to form compounds? Why do the sodium atom and the chlorine atom form sodium chloride? For many molecules, the constituent atoms have bonded according to the octet rule, which states that an atom tends to bond with other atoms until it has eight electrons in its outermost shell, thereby forming a stable electron configuration similar to that of the noble gases. However, this is not a hard and fast rule, more a “rule of thumb,” and as we suggested in the first chapter, there are more elements that can be exceptions to the rule than there are elements that follow the rule without exception. These “exceptional” elements include hydrogen, which can only have two valence electrons (achieving the configuration of helium); lithium and beryllium, which bond to attain two and four valence electrons, respectively; boron, which bonds to attain six valence electrons; and all elements in period 3 and below, which can expand the valence shell to include more than eight electrons by incorporating
Mnemonic
The
A simple way to remember all the exceptions is as follows:
• Hydrogen is excused from the octet rule because it doesn’t have enough “space” for eight electrons, it only has the one
• Lithium, beryllium, and boron are just lazy—they have enough room because they have both
• All the elements in period 3 and below have extra storage space in their attics, so they can hold more than eight electrons if they want to.
Another way to remember the exceptions (and one that’s even easier) is to remember the common elements that almost always abide by the octet rule: carbon, nitrogen, oxygen, fluorine, sodium, and magnesium. (We include the last two even though they lose—rather than gain—sufficient electrons to end up with a completed octet.)
