《有机化学》课程教学资源（参考书籍）Organic chemistry with biological applications, 3ed-Gengage Learning, John E. McMurry, Brooks Cole（2015）.pdf_P31-P35

P r o B l e m 1 . 1 Give the ground-state electron configuration for each of the following elements: (a) Oxygen (b) Phosphorus (c) Sulfur P r o B l e m 1 . 2 How many electrons does each of the following biological trace elements have in its outermost electron shell? (a) Magnesium (b) Cobalt (c) Selenium 1-4 Development of Chemical Bonding theory By the mid-1800s, the new science of chemistry was developing rapidly and chemists had begun to probe the forces holding atoms together in compounds. In 1858, August Kekulé and Archibald Couper independently proposed that, in all its compounds, carbon is tetravalent—it always forms four bonds when it joins other elements to form stable compounds. Furthermore, said Kekulé, carbon atoms can bond to one another to form extended chains of linked atoms. Shortly after the tetravalent nature of carbon was proposed, extensions to the Kekulé–Couper theory were made when the possibility of multiple bonding between atoms was suggested. Emil Erlenmeyer proposed a carbon–carbon triple bond for acetylene, and Alexander Crum Brown proposed a carbon– carbon double bond for ethylene. In 1865, Kekulé provided another major advance when he suggested that carbon chains can double back on themselves to form rings of atoms. Although Kekulé and Couper were correct in describing the tetravalent nature of carbon, chemistry was still viewed in a two-dimensional way until 1874. In that year, Jacobus van’t Hoff and Joseph Le Bel added a third dimension to our ideas about organic compounds. They proposed that the four bonds of carbon are not oriented randomly but have specific spatial directions. Van’t Hoff went even further and suggested that the four atoms to which carbon is bonded sit at the corners of a regular tetrahedron, with carbon in the center. A representation of a tetrahedral carbon atom is shown in FigUre 1.7. Note the conventions used to show three-dimensionality: solid lines represent bonds in the plane of the page, the heavy wedged line represents a bond coming out of the page toward the viewer, and the dashed line represents a bond receding back behind the page away from the viewer. These representations will be used throughout this text. H H H H Bond receding into page Bonds in plane of page Bond coming out of plane A tetrahedral carbon atom A regular tetrahedron C FigUre 1.7 a representation of van’t hoff’s tetrahedral carbon atom. The solid lines represent bonds in the plane of the paper, the heavy wedged line represents a bond coming out of the plane of the page, and the dashed line represents a bond going back behind the plane of the page. Unless otherwise noted, all content on this page is © Cengage Learning. 1-4 development of chemical Bonding theory 7 42912_01_Ch01_0001-0027h.indd 7 1/10/14 11:40 AM Copyright 2015 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it

8 CHAPTER 1 STRUCTURE AND BONDING Why.though,do atoms bond together,and how can bonds be described electronically?The why question is relatively easy to answer:atoms bond together because the compound that results is more stable and lower in energy than the separate atoms.Energy(usually as heat)is always released and flows out of the chemical system when a chemical bond forms.Conversely,energy must always be put into the system to break a chemical bond.Making bonds always releases energy,and breaking bonds always absorbs energy.The how question is more difficult.To answer it,we need to know more about the elec- tronic properties of atoms. We know through observation that eight electrons(an electron octet)in an atom's outermost shell,or valence she e the perioc e(2 -8);Ar (2 8+8) also knov emistry of a -group is gove dency to take on the electron co est noble gas nguration of the near etal r ex mple.ac eve a ga ing the s ingle s el rom t hal not urat o fill th n group 7A hel 1 ds like Na nion. :C The res -by an electrostatic attrac tion of unlike rges to the riodic table fo Look at methane.cH the main onstituent of nat for bonding in methane is not ionic because it would take too much ene carbon (1s2 2s2 2p2)to either pain or lose four electrons to achie ve a nobl electrons,but by sharing them.Such a shared-electron bond,first proposed in 1916 by G.N.Lewis,is called a covalent bond.The neutral collection ofatoms held together by covalent bonds is called a molecule. A simple way of indicating the covalent bonds in molecules is to use what are called Lewis structures,or electron-dot structures,in which the valence shell electrons of an atom are represented as dots.Thus,hydrogen has one dot senting its 1s electron,carbon has four dots (2s2p),oxygen has six dots -2p),and so on.A stable molecule results whenever a noble-gas configu ration is achieved for all the atoms eight dots(an octet)for main-group atoms drawn between atoms Elect H:N:H H:0:H H-C-H H-6-H -0- Methane Am Methan (CHa) (CH3OH) The number of covalent bonds an atom forms depends on how many addi- tional valence electrons it needs to reach a noble-gas configuration.Hydrogen

Why, though, do atoms bond together, and how can bonds be described electronically? The why question is relatively easy to answer: atoms bond together because the compound that results is more stable and lower in energy than the separate atoms. Energy (usually as heat) is always released and flows out of the chemical system when a chemical bond forms. Conversely, energy must always be put into the system to break a chemical bond. Making bonds always releases energy, and breaking bonds always absorbs energy. The how question is more difficult. To answer it, we need to know more about the electronic properties of atoms. We know through observation that eight electrons (an electron octet) in an atom’s outermost shell, or valence shell, impart special stability to the noblegas elements in group 8A of the periodic table: Ne (2 1 8); Ar (2 1 8 1 8); Kr (2 1 8 1 18 1 8). We also know that the chemistry of main-group elements is governed by their tendency to take on the electron configuration of the nearest noble gas. The alkali metals in group 1A, for example, achieve a noble-gas configuration by losing the single s electron from their valence shell to form a cation, while the halogens in group 7A achieve a noble-gas configuration by gaining a p electron to fill their valence shell and form an anion. The resultant ions are held together in compounds like Na1 Cl2 by an electrostatic attraction of unlike charges that we call an ionic bond. But how do elements closer to the middle of the periodic table form bonds? Look at methane, CH4, the main constituent of natural gas, for example. The bonding in methane is not ionic because it would take too much energy for carbon (1s2 2s2 2p2) to either gain or lose four electrons to achieve a noble-gas configuration. Instead, carbon bonds to other atoms, not by gaining or losing electrons, but by sharing them. Such a shared-electron bond, first proposed in 1916 by G. N. Lewis, is called a covalent bond. The neutral collection of atoms held together by covalent bonds is called a molecule. A simple way of indicating the covalent bonds in molecules is to use what are called Lewis structures, or electron-dot structures, in which the valenceshell electrons of an atom are represented as dots. Thus, hydrogen has one dot representing its 1s electron, carbon has four dots (2s2 2p2), oxygen has six dots (2s2 2p4), and so on. A stable molecule results whenever a noble-gas configuration is achieved for all the atoms—eight dots (an octet) for main-group atoms or two dots for hydrogen. Simpler still is the use of Kekulé structures, or linebond structures, in which two-electron covalent bonds are indicated as lines drawn between atoms. C HH H H H C H H N HH H O HH O H C HH H H N HH H H O Water (H2O) H H C H H Methane (CH4) Electron-dot structures (Lewis structures) Line-bond structures (Kekulé structures) Ammonia (NH3) Methanol (CH3OH) O H The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble-gas configuration. Hydrogen Unless otherwise noted, all content on this page is © Cengage Learning. 8 chapter 1 Structure and Bonding 42912_01_Ch01_0001-0027h.indd 8 1/10/14 11:40 AM Copyright 2015 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it

has one valence electron (1s) and needs one more to reach the helium configuration (1s2), so it forms one bond. Carbon has four valence electrons (2s2 2p2) and needs four more to reach the neon configuration (2s2 2p6), so it forms four bonds. Nitrogen has five valence electrons (2s2 2p3), needs three more, and forms three bonds; oxygen has six valence electrons (2s2 2p4), needs two more, and forms two bonds; and the halogens have seven valence electrons, need one more, and form one bond. One bond Four bonds Three bonds Two bonds One bond Br F Cl I H C N O Valence electrons that are not used for bonding are called lone-pair electrons, or nonbonding electrons. The nitrogen atom in ammonia (NH3), for instance, shares six valence electrons in three covalent bonds and has its remaining two valence electrons in a nonbonding lone pair. As a timesaving shorthand, nonbonding electrons are often omitted when drawing line-bond structures, but you still have to keep them in mind since they’re often crucial in chemical reactions. Nonbonding, lone-pair electrons N HH H or N HH or H N HH H Ammonia Predicting the number of Bonds Formed by atoms in molecules How many hydrogen atoms does phosphorus bond to in phosphine, PH?? S t r a t e g y Identify the periodic group of phosphorus, and tell from that how many electrons (bonds) are needed to make an octet. S o l u t i o n Phosphorus, like nitrogen, is in group 5A of the periodic table and has five valence electrons. It thus needs to share three more electrons to make an octet and therefore bonds to three hydrogen atoms, giving PH3. Drawing electron-Dot and line-Bond Structures Draw both electron-dot and line-bond structures for chloromethane, CH3Cl. S t r a t e g y Remember that a bond—that is, a pair of shared electrons—is represented as a line between atoms. W O R K E D E X A M P L E 1 . 1 W O R K E D E X A M P L E 1 . 2 Unless otherwise noted, all content on this page is © Cengage Learning. 1-4 development of chemical Bonding theory 9 42912_01_Ch01_0001-0027h.indd 9 1/10/14 11:40 AM Copyright 2015 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it

S o l u t i o n Hydrogen has one valence electron, carbon has four valence electrons, and chlorine has seven valence electrons. Thus, chloromethane is represented as CH Chloromethane H H CH H H Cl Cl P r o B l e m 1 . 3 Draw a molecule of chloroform, CHCl3, using solid, wedged, and dashed lines to show its tetrahedral geometry. P r o B l e m 1 . 4 Convert the adjacent representation of ethane, C2H6, into a conventional drawing that uses solid, wedged, and dashed lines to indicate tetrahedral geometry around each carbon (gray 5 C, ivory 5 H). P r o B l e m 1 . 5 What are likely formulas for the following substances? (a) CH?Cl2 (b) CH3SH? (c) CH3NH? P r o B l e m 1 . 6 Draw line-bond structures for the following substances, showing all nonbonding electrons: (a) CH3CH2OH, ethanol (b) H2S, hydrogen sulfide (c) CH3NH2, methylamine (d) N(CH3)3, trimethylamine P r o B l e m 1 . 7 Why can’t an organic molecule have the formula C2H7? 1-5 Describing Chemical Bonds: Valence Bond theory How does electron sharing lead to bonding between atoms? Two models have been developed to describe covalent bonding: valence bond theory and molecular orbital theory. Each model has its strengths and weaknesses, and chemists tend to use them interchangeably depending on the circumstances. Valence bond theory is the more easily visualized of the two, so most of the descriptions we’ll use in this book derive from that approach. According to valence bond theory, a covalent bond forms when two atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom. The electrons are now paired in the overlapping orbitals and are attracted to the nuclei of both atoms, thus bonding the atoms together. In the H2 molecule, for example, the H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals: + 2H H H 1s 1s H2 molecule H1 ) Ethane Unless otherwise noted, all content on this page is © Cengage Learning. 10 chapter 1 Structure and Bonding 42912_01_Ch01_0001-0027h.indd 10 1/10/14 11:40 AM Copyright 2015 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it

The overlapping orbitals in the H2 molecule have the elongated egg shape we might get by pressing two spheres together. If a plane were to pass through the middle of the bond, the intersection of the plane and the overlapping orbitals would be a circle. In other words, the H–H bond is cylindrically symmetrical, as shown in FigUre 1.8. Such bonds, which are formed by the head-on overlap of two atomic orbitals along a line drawn between the nuclei, are called sigma (s) bonds. During the bond-forming reaction 2 H· n H2, 436 kJ/mol (104 kcal/mol) of energy is released. Because the product H2 molecule has 436 kJ/mol less energy than the starting 2 H· atoms, the product is more stable than the reactants and we say that the H–H bond has a bond strength of 436 kJ/mol. In other words, we would have to put 436 kJ/mol of energy into the H–H bond to break the H2 molecule apart into H atoms (FigUre 1.9). [For convenience, we’ll generally give energies in both kilocalories (kcal) and the SI unit kilojoules (kJ): 1 kJ 5 0.2390 kcal; 1 kcal 5 4.184 kJ.] Two hydrogen atoms 2 H H2 H2 molecule 436 kJ/mol Released when bond forms Energy Absorbed when bond breaks How close are the two nuclei in the H2 molecule? If they are too close, they will repel each other because both are positively charged, yet if they are too far apart, they won’t be able to share the bonding electrons. Thus, there is an optimum distance between nuclei that leads to maximum stability (FigUre 1.10). Called the bond length, this distance is 74 pm in the H2 molecule. Every covalent bond has both a characteristic bond strength and bond length. HH (too close) Bond length 74 pm H H (too far) 0 + – H H Internuclear distance Energy Circular cross-section H H FigUre 1.8 Cylindrical symmetry of the h–h s bond. The intersection of a plane cutting through the s bond is a circle. FigUre 1.9 relative energy levels of two h atoms and the h2 molecule. The H2 molecule has 436 kJ/mol (104 kcal/mol) less energy than the two H atoms, so 436 kJ/mol of energy is released when the h–h bond forms. Conversely, 436 kJ/mol is absorbed when the h–h bond breaks. FigUre 1.10 a plot of energy versus internuclear distance for two hydrogen atoms. The distance between nuclei at the minimum energy point is the bond length. Unless otherwise noted, all content on this page is © Cengage Learning. 1-5 deScriBing chemical BondS: valence Bond theory 11 42912_01_Ch01_0001-0027h.indd 11 1/10/14 11:40 AM Copyright 2015 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it