《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

CHAPTER ONE Chemical Bonding PROBLEM 1.3 Which of the following ions possess a noble gas electron config- uration? (a)K (c)H (f)Ca2 SAMPLE SOLUTION (a)Potassium has atomic number 19, and so a potassium atom has 19 electrons. The ion k therefore. has 18 electrons the same as the noble gas argon. The electron configurations of k and Ar are the same 252p63523p Transfer of an electron from a sodium atom to a chlorine atom yields a sodium cation and a chloride anion, both of which have a noble gas electron configuration Na(g) ci(g) Sodium atom Chlorine atom Were we to simply add the ionization energy of sodium (496 kJ/ mol) and the electron affinity of chlorine(-349 kJ/mol), we would conclude that the overall process is endothermic with AH =+147 kJ/mol. The energy liberated by adding an electron to by the German physicist Wal. chlorine is insufficient to override the energy required to remove an electron from ter Kossel in 1916, in order sodium. This analysis, however, fails to consider the force of attraction between the to explain the ability of sub- oppositely charged ions Na and CI, which exceeds 500 kJ/mol and is more than suf- de to conduct an electric ficient to make the overall process exothermic. Attractive forces between oppositely charged particles are termed electrostatic, or coulombic, attractions and are what we mean by an ionic bond between two atoms PROBLEM 1. 4 What is the electron configuration of c of c Does either one of these ions have a noble gas(closed-shell)electron configuration? onic bonds are very common in inorganic compounds, but rare in organic ones. The ionization energy of carbon is too large and the electron affinity too small for car- bon to realistically form a C4+ or C4- ion. What kinds of bonds, then, link carbon to other elements in millions of organic compounds? Instead of losing or gaining electrons, carbon shares electrons with other elements (including other carbon atoms) to give wha are called covalent bonds 1.3 COVALENT BONDS The covalent, or shared electron pair, model of chemical bonding was first suggested Gilbert Newton Lewis (born by G. N. Lewis of the University of California in 1916. Lewis proposed that a sharing of two electrons by two hydrogen atoms permits each one to have a stable closed-shell eley, Califor- electron configuration analogous to helium nia, 1946)has been called he greatest American 984is H: H ue of the journal of chemi. cal Education contains fi Two hydrogen atoms. Hydrogen molecule: articles describing Lewis'life each with a single alent bonding by way of and contributions to chem. electron a shared electron pair Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

PROBLEM 1.3 Which of the following ions possess a noble gas electron config￾uration? (a) K (d) O (b) He (e) F (c) H (f) Ca2 SAMPLE SOLUTION (a) Potassium has atomic number 19, and so a potassium atom has 19 electrons. The ion K, therefore, has 18 electrons, the same as the noble gas argon. The electron configurations of K and Ar are the same: 1s 2 2s 2 2p6 3s 2 3p6 . Transfer of an electron from a sodium atom to a chlorine atom yields a sodium cation and a chloride anion, both of which have a noble gas electron configuration: Were we to simply add the ionization energy of sodium (496 kJ/mol) and the electron affinity of chlorine (349 kJ/mol), we would conclude that the overall process is endothermic with H° 147 kJ/mol. The energy liberated by adding an electron to chlorine is insufficient to override the energy required to remove an electron from sodium. This analysis, however, fails to consider the force of attraction between the oppositely charged ions Na and Cl– , which exceeds 500 kJ/mol and is more than suf- ficient to make the overall process exothermic. Attractive forces between oppositely charged particles are termed electrostatic, or coulombic, attractions and are what we mean by an ionic bond between two atoms. PROBLEM 1.4 What is the electron configuration of C? Of C? Does either one of these ions have a noble gas (closed-shell) electron configuration? Ionic bonds are very common in inorganic compounds, but rare in organic ones. The ionization energy of carbon is too large and the electron affinity too small for car￾bon to realistically form a C4 or C4 ion. What kinds of bonds, then, link carbon to other elements in millions of organic compounds? Instead of losing or gaining electrons, carbon shares electrons with other elements (including other carbon atoms) to give what are called covalent bonds. 1.3 COVALENT BONDS The covalent, or shared electron pair, model of chemical bonding was first suggested by G. N. Lewis of the University of California in 1916. Lewis proposed that a sharing of two electrons by two hydrogen atoms permits each one to have a stable closed-shell electron configuration analogous to helium. H Two hydrogen atoms, each with a single electron H Hydrogen molecule: covalent bonding by way of a shared electron pair H H Na(g) ±£ Sodium atom NaCl(g) Sodium chloride Cl(g) Chlorine atom  12 CHAPTER ONE Chemical Bonding Ionic bonding was proposed by the German physicist Wal￾ter Kossel in 1916, in order to explain the ability of sub￾stances such as sodium chlo￾ride to conduct an electric current. Gilbert Newton Lewis (born Weymouth, Massachusetts, 1875; died Berkeley, Califor￾nia, 1946) has been called the greatest American chemist. The January 1984 is￾sue of the Journal of Chemi￾cal Education contains five articles describing Lewis’ life and contributions to chem￾istry. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

1.3 Covalent Bonds Structural formulas of this type in which electrons are represented as dots are called Lewis structures The amount of energy required to dissociate a hydrogen molecule H2 to two sep arate hydrogen atoms is called its bond dissociation energy (or bond energy). For H it is quite large, being equal to 435 k/mol(104 kcal/mol). The main contributor to the strength of the covalent bond in H, is the increased binding force exerted on its two electrons. Each electron in H,"feels"the attractive force of two nuclei, rather than one as it would in an isolated hydrogen atom Covalent bonding in F2 gives each fluorine 8 electrons in its valence shell and a stable electron configuration equivalent to that of the noble gas neon: Two fluorine atoms. each with seven electrons in nding by way of ts valence shell electron pai PROBLEM 1.5 Hydrogen is bonded to fluorine in hydrogen fluoride by a cova- lent bond. Write a Lewis formula for hydrogen fluoride The Lewis model limits second-row elements(Li, Be, B, C, N, O, F, Ne) to a total of 8 electrons(shared plus unshared) in their valence shells. Hydrogen is limited to 2. Most of the elements that we'll encounter in this text obey the octet rule: in forming compounds they gain, lose, or share electrons to give a stable electron configuration characterized by eight valence electrons. When the octet rule is satisfied for carbon nitrogen, oxygen, and fluorine, they have an electron configuration analogous to the noble gas neon Now lets apply the Lewis model to the organic compounds methane and carbon tetrafluoride to write a Combine. C. and four h Lewis structure H: C: H for methane H Combine·C· and four Lewis structure F: C: F for carbon tetrafluoride Carbon has electrons in its valence shell in both methane and carbon tetrafluoride. By forming covalent bonds to four other atoms, carbon achieves a stable electron configu ration analogous to neon. Each covalent bond in methane and carbon tetrafluoride is quite strong--comparable to the bond between hydrogens in H2 in bond dissociation energy PROBLEM 1.6 Given the information that it has a carbon -carbon bond write a satisfactory Lewis structure for C2H(ethane) Representing a 2-electron covalent bond by a dash () the Lewis structures for hydrogen fluoride, fluorine, methane, and carbon tetrafluoride become Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

Structural formulas of this type in which electrons are represented as dots are called Lewis structures. The amount of energy required to dissociate a hydrogen molecule H2 to two sep￾arate hydrogen atoms is called its bond dissociation energy (or bond energy). For H2 it is quite large, being equal to 435 kJ/mol (104 kcal/mol). The main contributor to the strength of the covalent bond in H2 is the increased binding force exerted on its two electrons. Each electron in H2 “feels” the attractive force of two nuclei, rather than one as it would in an isolated hydrogen atom. Covalent bonding in F2 gives each fluorine 8 electrons in its valence shell and a stable electron configuration equivalent to that of the noble gas neon: PROBLEM 1.5 Hydrogen is bonded to fluorine in hydrogen fluoride by a cova￾lent bond. Write a Lewis formula for hydrogen fluoride. The Lewis model limits second-row elements (Li, Be, B, C, N, O, F, Ne) to a total of 8 electrons (shared plus unshared) in their valence shells. Hydrogen is limited to 2. Most of the elements that we’ll encounter in this text obey the octet rule: in forming compounds they gain, lose, or share electrons to give a stable electron configuration characterized by eight valence electrons. When the octet rule is satisfied for carbon, nitrogen, oxygen, and fluorine, they have an electron configuration analogous to the noble gas neon. Now let’s apply the Lewis model to the organic compounds methane and carbon tetrafluoride. Carbon has 8 electrons in its valence shell in both methane and carbon tetrafluoride. By forming covalent bonds to four other atoms, carbon achieves a stable electron configu￾ration analogous to neon. Each covalent bond in methane and carbon tetrafluoride is quite strong—comparable to the bond between hydrogens in H2 in bond dissociation energy. PROBLEM 1.6 Given the information that it has a carbon–carbon bond, write a satisfactory Lewis structure for C2H6 (ethane). Representing a 2-electron covalent bond by a dash (—), the Lewis structures for hydrogen fluoride, fluorine, methane, and carbon tetrafluoride become: Combine to write a Lewis structure for methane and fourC H CH H H H Combine to write a Lewis structure for carbon tetrafluoride and fourC F F F F F C Fluorine molecule: covalent bonding by way of a shared electron pair F F Two fluorine atoms, each with seven electrons in its valence shell F F 1.3 Covalent Bonds 13 Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

CHAPTER ONE Chemical Bonding F:H-C一H:F一C一F: H Methane Carbon tetrafluoride 1.4 DOUBLE BONDS AND TRIPLE BONDS Lewis's concept of shared electron pair bonds allows for 4-electron double bonds and 6-electron triple bonds. Carbon dioxide(CO2)has two carbon-oxygen double bonds and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide(HCN) has a carbon-nitrogen triple bor Hy drogen cyanide H:C∷:N H一 Multiple bonds are very common in organic chemistry. Ethylene( C2H4)contains a carbon-carbon double bond in its most stable lewis structure. and each carbon has a ipleted octet. The most stable Lewis structure for acetylene( C2H2) contains a car- bon-carbon triple bond. Here again, the octet rule is satisfied. H H H: C: C: H H一C≡C-H PROBLEM 1.7 Write the most stable Lewis structure for each of the following (a)Formaldehyde, CH,O. Both hydrogens are bonded to carbon. (A solution of hyde in water is sometimes used to Biological specimens. (b)Tetrafluoroethylene, C2F4. (The starting material for the preparation of Teflon. (c) Acrylonitrile, C3HaN. The atoms are connected in the order CCCN, and all hydrogens are bonded to carbon. (The starting material for the preparation of acrylic fibers such as Orlon and Acrilan) SAMPLE SoLUTION (a)Each hydrogen contributes 1 valence carbon contributes 4, and oxygen 6 for a total of 12 valence electrons told that both hydrogens are bonded to carbon. Since carbon forms four n its sta- ble compounds, join carbon and oxygen by a double bond. The partial structure so generated accounts for 8 of the 12 electrons. Add the remaining four electrons to oxygen as unshared pairs to complete the structure of formaldehyd H Partial structure showing Complete Lewis structure Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

1.4 DOUBLE BONDS AND TRIPLE BONDS Lewis’s concept of shared electron pair bonds allows for 4-electron double bonds and 6-electron triple bonds. Carbon dioxide (CO2) has two carbon–oxygen double bonds, and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon–nitrogen triple bond. Multiple bonds are very common in organic chemistry. Ethylene (C2H4) contains a carbon–carbon double bond in its most stable Lewis structure, and each carbon has a completed octet. The most stable Lewis structure for acetylene (C2H2) contains a car￾bon–carbon triple bond. Here again, the octet rule is satisfied. PROBLEM 1.7 Write the most stable Lewis structure for each of the following compounds: (a) Formaldehyde, CH2O. Both hydrogens are bonded to carbon. (A solution of formaldehyde in water is sometimes used to preserve biological specimens.) (b) Tetrafluoroethylene, C2F4. (The starting material for the preparation of Teflon.) (c) Acrylonitrile, C3H3N. The atoms are connected in the order CCCN, and all hydrogens are bonded to carbon. (The starting material for the preparation of acrylic fibers such as Orlon and Acrilan.) SAMPLE SOLUTION (a) Each hydrogen contributes 1 valence electron, carbon contributes 4, and oxygen 6 for a total of 12 valence electrons. We are told that both hydrogens are bonded to carbon. Since carbon forms four bonds in its sta￾ble compounds, join carbon and oxygen by a double bond. The partial structure so generated accounts for 8 of the 12 electrons. Add the remaining four electrons to oxygen as unshared pairs to complete the structure of formaldehyde. Partial structure showing covalent bonds O X C O X C H H ± ± Complete Lewis structure of formaldehyde H H ± ± orEthylene: C H H H H C CœC H H H H ± ± ± ± H H C orAcetylene: H±CPC±HC O orCarbon dioxide: O C OœCœO H N H orHydrogen cyanide: C ±CPN H±C±H H W W H Methane Carbon tetrafluoride F±C±F F W W F Hydrogen fluoride H±F Fluorine F±F 14 CHAPTER ONE Chemical Bonding Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

1.5 Polar Covalent Bonds and Electronegativity 1.5 POLAR COVALENT BONDS AND ELECTRONEGATIVITY Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect. If one atom has a greater tendency to attract electrons toward itself than the other we say the electron distribution is polarized, and the bond is referred to as a polar cova lent bond. Hydrogen fluoride, for example, has a polar covalent bond. Because fluorine attracts electrons more strongly than hydrogen, the electrons in the H-F bond are pulled toward fluorine, giving it a partial negative charge, and away from hydrogen giving it a partial positive charge. This polarization of electron density is represented in various ways 6H—F6- H一F CThe symbols and (The symbol +represents indicate partial positive the direction of polarization and partial negativ of electrons in the h-F bond) y of an atom to draw the electrons in a covalent bond toward itself is referred to electronegativity. An electronegative element attracts electrons: an electropositive one donates them. Electronegativity increases across a row in the peri- odic table. The most electronegative of the second-row elements is fluorine; the most electropositive is lithium Electronegativity decreases in going down a column. Fluorine is more electronegative than chlorine. The most commonly cited electronegativity scale was devised by Linus Pauling and is presented in Table 1. 2. PROBLEM 1.8 Examples of carbon-containing compounds include methane( CHa) chloromethane(CHaCi), and methyllithium(CH3 Li). In which one does carbon bear the greatest partial positive charge? The greatest partial negative charge? Technology, where he Centers of positive and negative charge that are separated from each other consti- in 1925. In addition to re tute a dipole. The dipole moment u of a molecule is equal to the charge e(either the pauling studied the structure between the centers of charge: the Nobel Prize in chemistry for that work in 1954. Paul efforts to limit the testing of uclear weapons. He was TABLE 1.2 Selected Values from the Pauling Electronegativity Scale ive won two Nobel prizes a woman. Can you name Group number Period H2u B C N 3.0 18 0.8 1.0 2.8 2.5 Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

1.5 POLAR COVALENT BONDS AND ELECTRONEGATIVITY Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect. If one atom has a greater tendency to attract electrons toward itself than the other, we say the electron distribution is polarized, and the bond is referred to as a polar cova￾lent bond. Hydrogen fluoride, for example, has a polar covalent bond. Because fluorine attracts electrons more strongly than hydrogen, the electrons in the H±F bond are pulled toward fluorine, giving it a partial negative charge, and away from hydrogen giving it a partial positive charge. This polarization of electron density is represented in various ways. The tendency of an atom to draw the electrons in a covalent bond toward itself is referred to as its electronegativity. An electronegative element attracts electrons; an electropositive one donates them. Electronegativity increases across a row in the peri￾odic table. The most electronegative of the second-row elements is fluorine; the most electropositive is lithium. Electronegativity decreases in going down a column. Fluorine is more electronegative than chlorine. The most commonly cited electronegativity scale was devised by Linus Pauling and is presented in Table 1.2. PROBLEM 1.8 Examples of carbon-containing compounds include methane (CH4), chloromethane (CH3Cl), and methyllithium (CH3Li). In which one does carbon bear the greatest partial positive charge? The greatest partial negative charge? Centers of positive and negative charge that are separated from each other consti￾tute a dipole. The dipole moment of a molecule is equal to the charge e (either the positive or the negative charge, since they must be equal) multiplied by the distance between the centers of charge:  e  d (The symbols  and  indicate partial positive and partial negative charge, respectively) H±F  H±F (The symbol represents the direction of polarization of electrons in the H±F bond) 1.5 Polar Covalent Bonds and Electronegativity 15 TABLE 1.2 Selected Values from the Pauling Electronegativity Scale Group number Period 1 2 3 4 5 I H 2.1 Li 1.0 Na 0.9 K 0.8 II Be 1.5 Mg 1.2 Ca 1.0 III B 2.0 Al 1.5 IV C 2.5 Si 1.8 V N 3.0 P 2.1 VI O 3.5 S 2.5 VII F 4.0 Cl 3.0 Br 2.8 I 2.5 Linus Pauling (1901–1994) was born in Portland, Ore￾gon and was educated at Oregon State University and at the California Institute of Technology, where he earned a Ph.D. in chemistry in 1925. In addition to re￾search in bonding theory, Pauling studied the structure of proteins and was awarded the Nobel Prize in chemistry for that work in 1954. Paul￾ing won a second Nobel Prize (the Peace Prize) for his efforts to limit the testing of nuclear weapons. He was one of only four scientists to have won two Nobel Prizes. The first double winner was a woman. Can you name her? Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website

CHAPTER ONE Chemical Bonding Because the charge on an electron is 4.80X 10 electrostatic units(esu) and tances within a molecule typically fall in the 10 cm range, molecular dipole are on the order of 10 8 esu- cm. In order to simplify the reporting of dipole this value of 10 esu' cm is defined as a debye, D. Thus the experimentally determined honor of Peter Debye, a dipole moment of hydrogen fluoride, 1.7 X 10esucm is stated as 1.7 Dutch scientist who did Table 1.3 lists the dipole moments of various bond types. For H-F, H-Cl, ortant work in many H-Br, and H-I these"bond dipoles"are really molecular dipole moments. A polar was awarded the Nobel Prize molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, bu can't have a shape that causes all the individual bond dipoles to cancel. We will have more to say about this in Section 1. ll after we have developed a feeling for the three- dimensional shapes of molecules The bond dipoles in Table 1.3 depend on the difference in electronegativity of the bonded atoms and on the bond distance. The polarity of a C-H bond is relatively low; substantially less than a C-O bond, for example. Don't lose sight of an even more important difference between a C-H bond and a C-O bond, and that is the direction of the dipole moment. In a C-H bond the electrons are drawn away from H, toward C In a C-o bond, electrons are drawn from C toward O. As we'll see in later chap ters, the kinds of reactions that a substance undergoes can often be related to the size and direction of key bond dipoles 1.6 FORMAL CHARGE Lewis structures frequently contain atoms that bear a positive or negative charge. molecule as a whole is neutral, the sum of its positive charges must equal the sum of its negative charges. An example is nitric acid, HNO As written, the structural formula for nitric acid depicts different bonding patterns for its three oxygens. One oxygen is doubly bonded to nitrogen, another is singly bonded TABLE 1.3 Selected Bond Dipole Moments Bond* Dipole moment, D Bond* Dipole moment, D H—F C-F H—C H—C C≡N 3.6 irection of the dipole moment is toward the more electronegative atom In the listed examples en and carbon are the positive ends of the dipoles. Carbon is the negative end of the dipole ted with the C-H bond Back Forward Main MenuToc Study Guide ToC Student o MHHE Website

Because the charge on an electron is 4.80  1010 electrostatic units (esu) and the dis￾tances within a molecule typically fall in the 108 cm range, molecular dipole moments are on the order of 1018 esu·cm. In order to simplify the reporting of dipole moments this value of 1018 esu cm is defined as a debye, D. Thus the experimentally determined dipole moment of hydrogen fluoride, 1.7  1018 esu cm is stated as 1.7 D. Table 1.3 lists the dipole moments of various bond types. For H±F, H±Cl, H±Br, and H±I these “bond dipoles” are really molecular dipole moments. A polar molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, but can’t have a shape that causes all the individual bond dipoles to cancel. We will have more to say about this in Section 1.11 after we have developed a feeling for the three￾dimensional shapes of molecules. The bond dipoles in Table 1.3 depend on the difference in electronegativity of the bonded atoms and on the bond distance. The polarity of a C±H bond is relatively low; substantially less than a C±O bond, for example. Don’t lose sight of an even more important difference between a C±H bond and a C±O bond, and that is the direction of the dipole moment. In a C±H bond the electrons are drawn away from H, toward C. In a C±O bond, electrons are drawn from C toward O. As we’ll see in later chap￾ters, the kinds of reactions that a substance undergoes can often be related to the size and direction of key bond dipoles. 1.6 FORMAL CHARGE Lewis structures frequently contain atoms that bear a positive or negative charge. If the molecule as a whole is neutral, the sum of its positive charges must equal the sum of its negative charges. An example is nitric acid, HNO3: As written, the structural formula for nitric acid depicts different bonding patterns for its three oxygens. One oxygen is doubly bonded to nitrogen, another is singly bonded H±O±N O O   œ ± 16 CHAPTER ONE Chemical Bonding TABLE 1.3 Selected Bond Dipole Moments Bond* H±F H±Cl H±Br H±I H±C H±N H±O Dipole moment, D 1.7 1.1 0.8 0.4 0.3 1.3 1.5 Bond* C±F C±O C±N CœO CœN CPN Dipole moment, D 1.4 0.7 0.4 2.4 1.4 3.6 *The direction of the dipole moment is toward the more electronegative atom. In the listed examples hydrogen and carbon are the positive ends of the dipoles. Carbon is the negative end of the dipole associated with the C±H bond. The debye unit is named in honor of Peter Debye, a Dutch scientist who did im￾portant work in many areas of chemistry and physics and was awarded the Nobel Prize in chemistry in 1936. Back Forward Main Menu TOC Study Guide TOC Student OLC MHHE Website