《有机化学》课程教学资源（教材文献，英文版）CHAPTER 1 CHEMICAL BONDING

c3● CHAPTER 1 CHEMICAL BONDING Sa 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 formula 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 principles 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 hroughout your study of organic chemistry. 1.1 ATOMS, ELECTRONS, AND ORBITALS Before discussing bonding principles, lets first 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, which 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 e properties as well. Two years later Erwin Schrodinger took the next step and cal- culated the energy of an electron in a hydrogen atom by using equations that treated the electron as if it were a wave. Instead of a single energy, Schrodinger obtained a series of energy levels, each of which corresponded to a different mathematical description of the electron wave. These mathematical descriptions are called 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. Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

7 CHAPTER 1 CHEMICAL BONDING Structure* 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 formula 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 principles 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. 1.1 ATOMS, ELECTRONS, AND ORBITALS Before discussing bonding principles, let’s first 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, which 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 Schrödinger took the next step and cal￾culated the energy of an electron in a hydrogen atom by using equations that treated the electron as if it were a wave. Instead of a single energy, Schrödinger obtained a series of energy levels, each of which corresponded to a different mathematical description of the electron wave. These mathematical descriptions are called 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. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

CHAPTER ONE Chemical Bonding elect According to the Heisenberg uncertainty principle, we can't tell exactly where an on is, but we can tell where it is most likely to be. The probability of finding an electron at a particular spot relative to an atom's nucleus is given by the square of the wave function( -)at that point. Figure 1. I illustrates the probability of finding an elec- tron at various 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 finding an elec- tron at a particular point is greatest near the nucleus, and decreases with increasing dis- tance from the nucleus but never becomes zero. We commonly describe Figure 1.1 as FIGURE 1.1 Probability dis- an"electron cloud to call attention to the spread-out nature of the electron probability tribution(4 )for an electron Be careful, though. The"electron cloud "of a hydrogen atom, although drawn as a col in a 1s orbital lection of many dots, represents only one electron Wave functions are also called orbitals. For convenience. chemists use the term orbital"in several different ways. a drawing such as Figure 1. I 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 the price to be paid or simplicity of expression Orbitals are described by specifying their size, shape, and directional properties Spherically symmetrical ones such as shown in Figure 1. I are called s orbitals. The let ter s is preceded by the principal quantum number n(n= 1, 2, 3, etc. )which speci fies the shell and is related to the energy of the orbital. An electron in a ls orbital is likely to be found closer to the nucleus, is lower in energy, and is more strongly held han an electron in a 2s orbital Regions of a single orbital may be separated by nodal surfaces where the proba bility of finding an electron is zero. A ls orbital has no nodes; a 2s orbital has one. A Is and a 2s orbital are shown in cross section in Figure 1. 2. The 2s wave function changes ign on passing through the nodal surface as indicated by the plus (+) and minus ( signs in Figure 1. 2. Do not confuse these signs with electric charges--they have noth ing to do with electron or nuclear charge. Also, be aware that our"orbital"drawings really representations of v-(which must be a positive number), whereas and refer to the sign of the wave function() itself. These customs may seem confusing at first but turn out not to complicate things in practice. Indeed, most of the time we wont Nucleus FIGURE 1.2 Cross 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 spherical node where the wave function changes Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

According to the Heisenberg uncertainty principle, we can’t tell exactly where an electron is, but we can tell where it is most likely to be. The probability of finding an electron at a particular spot relative to an atom’s nucleus is given by the square of the wave function (2 ) at that point. Figure 1.1 illustrates the probability of finding an elec￾tron at various 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 finding an elec￾tron at a particular point is greatest near the nucleus, and decreases with increasing dis￾tance from the nucleus but never becomes 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 called 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 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.) which speci- fies 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 proba￾bility of finding 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 confuse these signs with electric charges—they 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 confusing at first but turn out not to complicate things in practice. Indeed, most of the time we won’t 8 CHAPTER ONE Chemical Bonding x z y FIGURE 1.1 Probability dis￾tribution (2 ) for an electron in a 1s orbital. Node (a) (b) Nucleus y x
Nucleus x y
FIGURE 1.2 Cross 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 spherical node where the wave function changes sign. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

1.1 Atoms, Electrons, and Orbitals even include and -signs of wave functions in our drawings but only when they are necessary for understanding a particular concept Instead of probability distributions, it is more common to represent orbitals by their boundary surfaces, as shown in Figure 1.3 for the ls and 2s orbitals. The boundary sur face encloses the region where the probability of finding an electron is high--on the order of 90-95%. Like the probability distribution plot from which it is derived, a pic ture of a boundary surface is usually described as a drawing of an orbital single electron of hydrogen occupies a Is orbital, as do the two electrons of helium. The respective electron configurations are described as s Helium: 152 In addition to being negatively charged, electrons possess the property of spin. The spin quantum number of an electron can have a value of either +2 or -2 According to the Pauli exclusion principle, two electrons may occupy the same orbital only when they have opposite, or"paired, " spins. For this reason, no orbital can contain more than two electrons. Since two electrons fill the is orbital. the third electron in lithium (z=3)must occupy an orbital of higher energy. After 1s, the next higher energy orbital is 2s. The third electron in lithium therefore occupies the 2s orbital, and the electron configuration of lithium is Li case of hydrogen and helium). Hydrogen and helium are first-row elements; lithium the inside back presenteyof The period (or row) of the periodic table in which an element appears corresponds to the principal quantum number of the highest numbered occupied orbital (n= 1 in the the elements (n= 2)is a second-row element. with beryllium(Z =4), the 2s level becomes filled, and the next orbitals to be occupied in it and the remaining second-row elements are the 2po 2py, and 2p2 orbitals These orbitals, portrayed in Figure 1. 4, have a boundary surface that is described as"dumbbell-shaped. "Each orbital consists of two"lobes, "that is spheres that touch each other along a nodal plane passing through the nucleus. The 2p x 2py, and 2p, orbitals are equal in energy and mutually perpendicular. The electron configurations of the first 12 elements, hydrogen through magnesium, are given in Table 1. 1. In filling the 2p orbitals, notice that each is singly occupied before any one is doubly occupied. This is a general principle for orbitals of equal energy kno FIGURE 1.3 Bound urfaces of a 1s orbital and a 2s orbital. The boundary surfaces enclose the volume where there is a 90-95% probability of finding an electron Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

even include  and  signs of wave functions in our drawings but only when they are necessary for understanding a particular concept. Instead of probability distributions, it is more common to represent orbitals by their boundary surfaces, as shown in Figure 1.3 for the 1s and 2s orbitals. The boundary sur￾face encloses the region where the probability of finding an electron is high—on the order of 90–95%. Like the probability distribution plot from which it is derived, a pic￾ture of a boundary surface is usually described as a drawing of an orbital. A hydrogen atom (Z 1) has one electron; a helium atom (Z 2) has two. The single electron of hydrogen occupies a 1s orbital, as do the two electrons of helium. The respective electron configurations are described as: Hydrogen: 1s 1 Helium: 1s 2 In addition to being negatively charged, electrons possess the property of spin. The spin quantum number of an electron can have a value of either 1 2 or 1 2. According to the Pauli exclusion principle, two electrons may occupy the same orbital only when they have opposite, or “paired,” spins. For this reason, no orbital can contain more than two electrons. Since two electrons fill the 1s orbital, the third electron in lithium (Z 3) must occupy an orbital of higher energy. After 1s, the next higher energy orbital is 2s. The third electron in lithium therefore occupies the 2s orbital, and the electron configuration of lithium is Lithium: 1s 2 2s 1 The period (or row) of the periodic table in which an element appears corresponds to the principal quantum number of the highest numbered occupied orbital (n 1 in the case of hydrogen and helium). Hydrogen and helium are first-row elements; lithium (n 2) is a second-row element. With beryllium (Z 4), the 2s level becomes filled, and the next orbitals to be occupied in it and the remaining second-row elements are the 2px, 2py, and 2pz orbitals. These orbitals, portrayed in Figure 1.4, have a boundary surface that is usually described as “dumbbell-shaped.” Each orbital consists of two “lobes,” that is, slightly flattened spheres that touch each other along a nodal plane passing through the nucleus. The 2px, 2py, and 2pz orbitals are equal in energy and mutually perpendicular. The electron configurations of the first 12 elements, hydrogen through magnesium, are given in Table 1.1. In filling the 2p orbitals, notice that each is singly occupied before any one is doubly occupied. This is a general principle for orbitals of equal energy known 1.1 Atoms, Electrons, and Orbitals 9 z z x y y x 1s 2s FIGURE 1.3 Boundary surfaces of a 1s orbital and a 2s orbital. The boundary surfaces enclose the volume where there is a 90–95% probability of finding an electron. A complete periodic table of the elements is presented on the inside back cover. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

CHAPTER ONE Chemical Bonding 2p FIGURE 1.4 Boundary surfaces of the 2p s. The wave function changes sign at the ne is a nodal surface for the 2p, orbit and the xy-plane is a nodal surface for the 2p, orbital as Hunds rule. Of particular importance in Table 1. are hydrogen, carbon, nitrogen, and oxygen. Countless organic compounds contain nitrogen, oxygen, or both in addition to car bon, the essential element of organic chemistry. Most of them also contain hydrogen It is often convenient to speak of the valence electrons of an atom. These are the outermost electrons, the ones most likely to be involved in chemical bonding and reac tions. For second-row elements these are the 2s and 2p electrons. Because four orbitals (25, 2pr 2p, 2p,) are involved, the um number of electrons in the valence shell of any second-row element is 8. Neon, with all its 2s and 2p orbitals doubly occupied, has eight valence electrons and completes the second row of the periodic table Answers to all problems that PROBLEM 1.1 How many valence electrons does carbon have? appear within the body of a cs 2. A brief discussion of Once the 2s and 2p orbitals are filled, the next level is the 3s, followed by the 3po 3p nd 3p, orbitals. Electrons in these orbitals are farther from the nucleus than those in the ow to do problems of the same type are offered in the 2s and 2p orbitals and are of higher energy Study Guid TABLE 1.1 Electron Configurations of the First Twelve Elements of the Periodic Table Number of electrons in indicated orbital Atomi Element number z 1s 25 批m oron 12345678 Fluorine Neon 10 122222222222 1222222222 Sodium 22222 222 2 Magnesium 12 Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

as Hund’s rule. Of particular importance in Table 1.1 are hydrogen, carbon, nitrogen, and oxygen. Countless organic compounds contain nitrogen, oxygen, or both in addition to car￾bon, the essential element of organic chemistry. Most of them also contain hydrogen. It is often convenient to speak of the valence electrons of an atom. These are the outermost electrons, the ones most likely to be involved in chemical bonding and reac￾tions. For second-row elements these are the 2s and 2p electrons. Because four orbitals (2s, 2px, 2py, 2pz) are involved, the maximum number of electrons in the valence shell of any second-row element is 8. Neon, with all its 2s and 2p orbitals doubly occupied, has eight valence electrons and completes the second row of the periodic table. PROBLEM 1.1 How many valence electrons does carbon have? Once the 2s and 2p orbitals are filled, the next level is the 3s, followed by the 3px, 3py, and 3pz orbitals. Electrons in these orbitals are farther from the nucleus than those in the 2s and 2p orbitals and are of higher energy. 10 CHAPTER ONE Chemical Bonding x xx z y yy zz 2px 2pz 2py FIGURE 1.4 Boundary surfaces of the 2p orbitals. The wave function changes sign at the nucleus. The yz-plane is a nodal surface for the 2px orbital. The probability of finding a 2px electron in the yz-plane is zero. Analogously, the xz-plane is a nodal surface for the 2py orbital, and the xy-plane is a nodal surface for the 2pz orbital. TABLE 1.1 Electron Configurations of the First Twelve Elements of the Periodic Table Number of electrons in indicated orbital Element Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Atomic number Z 1 2 3 4 5 6 7 8 9 10 11 12 1s 1 2 2 2 2 2 2 2 2 2 2 2 2s 1 2 2 2 2 2 2 2 2 2 2px 1 1 1 2 2 2 2 2 2py 1 1 1 2 2 2 2 2pz 1 1 1 2 2 2 3s 1 2 Answers to all problems that appear within the body of a chapter are found in Appen￾dix 2. A brief discussion of the problem and advice on how to do problems of the same type are offered in the Study Guide. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

nic Bonds PROBLEM 1.2 Referring to the periodic table as needed, write electron config urations for all the elements in the third period SAMPLE SoLUTION The third period begins with sodium and ends with argon ompanied by a sample so The atomic number z of sodium is 11. and so a sodium atom has 11 electro the other parts of the prob. The maximum number of electrons in the 1s, 25, and 2p orbitals is ten, and so the und in Appendix 2, eleventh electron of sodium occupies a 3s orbital. The electron configuration of and detailed solutions are sodium is 15252px 22py 22p 23s ed in the Study Neon, in the second period, and argon, in the third, possess eight electrons in their valence shell; they are said to have a complete octet of electrons. Helium, neon, and argon belong to the class of elements known as noble gases or rare gases. The noble gases are characterized by an extremely stable"closed-shell"electron configuration and are very unreactive. 1.2 IONIC BONDS Atoms combine with one another to give compounds having properties different from the atoms they contain. The attractive force between atoms in a compound is a chemi cal bond. One type of chemical bond, called an ionic bond, is the force of attraction between oppositely charged species (ions)(Figure 1.5). lons that are positively charged are referred to as cations; those that are negatively charged are anions. FIGURE 1.5 An Whether an element is the source of the cation or anion in an ionic bond depends is the force of el on several factors, for which the periodic table can serve as a guide. In forming ionic attraction between compounds, elements at the left of the periodic table typically lose electrons, forming a sitely, charged ions, illus- cation that has the same electron configuration as the nearest noble gas. Loss of an elec tron from sodium, for example, gives the species Na, which has the same electron con- solid sodium chloride, each sodium ion is surrounded by six chloride ions and vice Sodium atom Electron A large amount of energy, called the ionization energy, must be added to any atom The sI (Systeme Int ternational in order to dislodge one of its electrons. The ionization energy of sodium, for example, d'" Unites) unit of energy is is 496 kJ/mol (119 kcal/mol). Processes that absorb energy are said to be endothermic. the joule o). An older unit is Compared with other elements, sodium and its relatives in group IA have relatively low ganic chemists still express ionization energies. In general, ionization energy increases across a row in the periodic ergy changes in units of table Elements at the right of the periodic table tend to gain electrons to reach the elec kcalmol =4.184 kJ/mol) tron configuration of the next higher noble gas. Adding an electron to chlorine, for exam- ple, gives the anion CI, which has the same closed-shell electron configuration as the noble gas argon ci(g) Chlorine atom Electron Chloride ion Energy is released when a chlorine atom captures an electron Energy-releasing reactions are described as exothermic, and the energy change for an exothermic process has a negative sign. The energy change for addition of an electron to an atom is referred to as its electron affinity and is -349 kJmol (-83 4 kcal/mol) for chlorine Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

PROBLEM 1.2 Referring to the periodic table as needed, write electron config￾urations for all the elements in the third period. SAMPLE SOLUTION The third period begins with sodium and ends with argon. The atomic number Z of sodium is 11, and so a sodium atom has 11 electrons. The maximum number of electrons in the 1s, 2s, and 2p orbitals is ten, and so the eleventh electron of sodium occupies a 3s orbital. The electron configuration of sodium is 1s 2 2s 2 2px 2 2py 2 2pz 2 3s 1 . Neon, in the second period, and argon, in the third, possess eight electrons in their valence shell; they are said to have a complete octet of electrons. Helium, neon, and argon belong to the class of elements known as noble gases or rare gases. The noble gases are characterized by an extremely stable “closed-shell” electron configuration and are very unreactive. 1.2 IONIC BONDS Atoms combine with one another to give compounds having properties different from the atoms they contain. The attractive force between atoms in a compound is a chemi￾cal bond. One type of chemical bond, called an ionic bond, is the force of attraction between oppositely charged species (ions) (Figure 1.5). Ions that are positively charged are referred to as cations; those that are negatively charged are anions. Whether an element is the source of the cation or anion in an ionic bond depends on several factors, for which the periodic table can serve as a guide. In forming ionic compounds, elements at the left of the periodic table typically lose electrons, forming a cation that has the same electron configuration as the nearest noble gas. Loss of an elec￾tron from sodium, for example, gives the species Na, which has the same electron con- figuration as neon. A large amount of energy, called the ionization energy, must be added to any atom in order to dislodge one of its electrons. The ionization energy of sodium, for example, is 496 kJ/mol (119 kcal/mol). Processes that absorb energy are said to be endothermic. Compared with other elements, sodium and its relatives in group IA have relatively low ionization energies. In general, ionization energy increases across a row in the periodic table. Elements at the right of the periodic table tend to gain electrons to reach the elec￾tron configuration of the next higher noble gas. Adding an electron to chlorine, for exam￾ple, gives the anion Cl, which has the same closed-shell electron configuration as the noble gas argon. Energy is released when a chlorine atom captures an electron. Energy-releasing reactions are described as exothermic, and the energy change for an exothermic process has a negative sign. The energy change for addition of an electron to an atom is referred to as its electron affinity and is 349 kJ/mol (83.4 kcal/mol) for chlorine. Cl(g) ±£ Chlorine atom 1s 2 2s 2 2p6 3s 2 3p5 Cl(g) Chloride ion 1s 2 2s 2 2p6 3s 2 3p6 e Electron  Na(g) ±£ Sodium atom 1s 2 2s 2 2p6 3s 1 [The (g) indicates that the species is present in the gas phase.] Na(g) Sodium ion 1s 2 2s 2 2p6 e Electron  1.2 Ionic Bonds 11
FIGURE 1.5 An ionic bond is the force of electrostatic attraction between oppo￾sitely charged ions, illus￾trated in this case by Na (red) and Cl (green). In solid sodium chloride, each sodium ion is surrounded by six chloride ions and vice versa in a crystal lattice. In-chapter problems that contain multiple parts are ac￾companied by a sample solu￾tion to part (a). Answers to the other parts of the prob￾lem are found in Appendix 2, and detailed solutions are presented in the Study Guide. The SI (Système International d’Unites) unit of energy is the joule (J). An older unit is the calorie (cal). Most or￾ganic chemists still express energy changes in units of kilocalories per mole (1 kcal/mol 4.184 kJ/mol). Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website