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1、7 CHAPTER 1 CHEMICAL BONDING S tructure* is the key to everything in chemistry. The properties of a substance depend on the atoms it contains and the way the atoms are connected. What is less obvious, but very powerful, is the idea that someone who is trained in chemistry can look at a structural fo
2、rmula of a substance and tell you a lot about its properties. This chapter begins your training toward understanding the relationship between struc- ture and properties in organic compounds. It reviews some fundamental principles of molecular structure and chemical bonding. By applying these princip
3、les you will learn to recognize the structural patterns that are more stable than others and develop skills in communicating chemical information by way of structural formulas that will be used throughout your study of organic chemistry. 1.1ATOMS, ELECTRONS, AND ORBITALS Before discussing bonding pr
4、inciples, lets fi rst review some fundamental relationships between atoms and electrons. Each element is characterized by a unique atomic number Z, which is equal to the number of protons in its nucleus. A neutral atom has equal num- bers of protons, which are positively charged, and electrons, whic
5、h are negatively charged. Electrons were believed to be particles from the time of their discovery in 1897 until 1924, when the French physicist Louis de Broglie suggested that they have wave- like properties as well. Two years later Erwin Schrdinger took the next step and cal- culated the energy of
6、 an electron in a hydrogen atom by using equations that treated the electron as if it were a wave. Instead of a single energy, Schrdinger obtained a series of energy levels, each of which corresponded to a different mathematical description of the electron wave. These mathematical descriptions are c
7、alled wave functions and are symbolized by the Greek letter ? (psi). *A glossary of important terms may be found immediately before the index at the back of the book. BackForwardMain MenuTOCStudy Guide TOCStudent OLCMHHE Website According to the Heisenberg uncertainty principle, we cant tell exactly
8、 where an electron is, but we can tell where it is most likely to be. The probability of fi nding an electron at a particular spot relative to an atoms nucleus is given by the square of the wave function (? 2 ) at that point. Figure 1.1 illustrates the probability of fi nding an elec- tron at variou
9、s points in the lowest energy (most stable) state of a hydrogen atom. The darker the color in a region, the higher the probability. The probability of fi nding an elec- tron at a particular point is greatest near the nucleus, and decreases with increasing dis- tance from the nucleus but never become
10、s zero. We commonly describe Figure 1.1 as an “electron cloud” to call attention to the spread-out nature of the electron probability. Be careful, though. The “electron cloud” of a hydrogen atom, although drawn as a col- lection of many dots, represents only one electron. Wave functions are also cal
11、led orbitals. For convenience, chemists use the term “orbital” in several different ways. A drawing such as Figure 1.1 is often said to repre- sent an orbital. We will see other kinds of drawings in this chapter, use the word “orbital” to describe them too, and accept some imprecision in language as
12、 the price to be paid for simplicity of expression. Orbitals are described by specifying their size, shape, and directional properties. Spherically symmetrical ones such as shown in Figure 1.1 are called s orbitals. The let- ter s is preceded by the principal quantum number n (n ? 1, 2, 3, etc.) whi
13、ch speci- fi es the shell and is related to the energy of the orbital. An electron in a 1s orbital is likely to be found closer to the nucleus, is lower in energy, and is more strongly held than an electron in a 2s orbital. Regions of a single orbital may be separated by nodal surfaces where the pro
14、ba- bility of fi nding an electron is zero. A 1s orbital has no nodes; a 2s orbital has one. A 1s and a 2s orbital are shown in cross section in Figure 1.2. The 2s wave function changes sign on passing through the nodal surface as indicated by the plus (?) and minus (?) signs in Figure 1.2. Do not c
15、onfuse these signs with electric chargesthey have noth- ing to do with electron or nuclear charge. Also, be aware that our “orbital” drawings are really representations of ? 2 (which must be a positive number), whereas ? and ? refer to the sign of the wave function (?) itself. These customs may seem
16、 confusing at fi rst but turn out not to complicate things in practice. Indeed, most of the time we wont 8CHAPTER ONEChemical Bonding x z y FIGURE 1.1Probability dis- tribution (?2) for an electron in a 1s orbital. Node (a)(b) Nucleus y x ? ? Nucleus x y ? FIGURE 1.2Cross sections of (a) a 1s orbital and (b) a 2s orbital. The wave function has the same sign over the entire 1s orbital. It is arbitrarily shown as ?, but could just as well have been designated as ?. The 2s orbital has a spheric
