A model of the structure of diamond,one form of pure Covalent Bonding and corners of a tetrahedron Shapes of Molecules Fullerenes. Outline 1.1Electronic Structure of Atoms 1.2 Lewis Model of Bonding HOW TO Quickly Figure Out Formal Charge HOW TO om Condensed 1.3Functional Groups 1.4 Bond Angles and Shapes of Molecules 1.5Polar and Nonpolar Molecules 1.6 Quantum or Wave Mechanics 1.7 A Combined Valence Bond and Molecular Orbital Theory Approach to Covalent Bonding HOW TO 8cccashebrdzaionand toms 1.8Resonance HOW TO Draw Curved Arrows and Push Electrons in Creating Contributing Structures 1.9 Molecular Orbitals for Delocalized Systems Go to www.cen Accordina to the simplest definition. is the study of the compounds ofarbn-Papornom pounds consist of carbon and only a few other elements-chiefly,hydrogen,oxygen, lectures for this chanter Unessisotelartnhn3
Unless otherwise noted all art on this page © Cengage Learning 2013 1 © Cengage Learning/Charles D. Winters A model of the structure of diamond, one form of pure carbon. Each carbon is bonded to four other carbons at the corners of a tetrahedron. Inset: a model of fullerene (C60). See “MCAT Practice: Fullerenes.” 1 Covalent Bonding and Shapes of Molecules Outline 1.1 Electronic Structure of Atoms 1.2 Lewis Model of Bonding How To Quickly Figure Out Formal Charge How To Draw Lewis Structures from Condensed Structural Formulas 1.3 Functional Groups 1.4 Bond Angles and Shapes of Molecules 1.5 Polar and Nonpolar Molecules 1.6 Quantum or Wave Mechanics 1.7 A Combined Valence Bond and Molecular Orbital Theory Approach to Covalent Bonding How To Quickly Recognize the Hybridization and Geometry of Atoms 1.8 Resonance How To Draw Curved Arrows and Push Electrons in Creating Contributing Structures 1.9 Molecular Orbitals for Delocalized Systems 1.10 Bond Lengths and Bond Strengths in Alkanes, Alkenes, and Alkynes Go to www.cengage.com/chemistry/ brown/organic7e and click Access Student Materials to view video lectures for this chapter. According to the simplest definition, organic chemistry is the study of the compounds of carbon. Perhaps its most remarkable feature is that most organic compounds consist of carbon and only a few other elements—chiefly, hydrogen, oxygen, Copyright 2013 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
Chapter t Bonding and and nitrogen.Chemists have discovered or made well over 10 million compounds Shapes of Molecules composed of carbon and ese thre ther elements.Organi compoun are a etics:our nte our glues and adhesives:in our fuels and lubricants:and,of course.in our bodies and the bodies of all living things. Let us review how the ele sof C.H,O and combine by sharing elec tonphsnital t,you s:h followrequire your use this knowledge luenty 1.1 Electronic Structure of Atoms -1010m neutrons nd Figure 1.1 cleus of an atom has a diameter of 10-14 to 10- 尚p red electrons.The nu- is in its small dense nucleus. probability rof approximately 10-m(Figure 1.1) ed.o ospace rel tive to the ru values.These shells occur only at quantized energies in which three important effects of space ctrostatic attraction that draws the lectrons to esecond is the preading the electron density awa from the nuclei.Delocalization describes the spreading of electron density over a larger volume of space. numbers 1,2, d.18 ele rons:the fourth 32 ele s:and so on (Table 1.1).Electr nearest to the positively charged nucleus and are held most strongly by it;these elec trons are lowest in energy.Electrons higher-numbered shells are farther from the ely ch us and are eld 1 ss,p,d,and f,and within Orbital hes itals (Table 1).Ar space that can hold two electrons and has a specific quantized ener.The first shel second shell contains ones orbital and states in lap.As a point of reference,to discuss the 2p orthogonal orbitals,we consider them to be directed along the x-.yand =-axes and give them designations,2p2p and Table 1.1 Distribution of Electrons in Shells Number of Relative Energie 1 8 lower 2
Unless otherwise noted all art on this page © Cengage Learning 2013 Chapter 1 Covalent Bonding and Shapes of Molecules 2 and nitrogen. Chemists have discovered or made well over 10 million compounds composed of carbon and these three other elements. Organic compounds are all around us—in our foods, flavors, and fragrances; in our medicines, toiletries, and cosmetics; in our plastics, films, fibers, and resins; in our paints and varnishes; in our glues and adhesives; in our fuels and lubricants; and, of course, in our bodies and the bodies of all living things. Let us review how the elements of C, H, O, and N combine by sharing electron pairs to form bonds, and ultimately molecules. No doubt, you have encountered much of this initial material in previous chemistry courses; however, the chapters that follow require your ability to use this knowledge fluently. 1.1 Electronic Structure of Atoms An atom contains a small, dense nucleus made of neutrons and positively charged protons. Most of the mass of an atom is contained in its nucleus. The nucleus is surrounded by an extranuclear space containing negatively charged electrons. The nucleus of an atom has a diameter of 10214 to 10215 meters (m). The electrons occupy a much larger volume with a diameter of approximately 10210 m (Figure 1.1). Shells define the probability of finding an electron in various regions of space relative to the nucleus. The energy of electrons in the shells is quantized. Quantization means that only specific values of energy are possible, rather than a continuum of values. These shells occur only at quantized energies in which three important effects balance each other. The first is the electrostatic attraction that draws the electrons toward the nucleus; the second is the electrostatic repulsion between the electrons; and the third is the wavelike nature of an electron that prefers to be delocalized, thereby spreading the electron density away from the nuclei. Delocalization describes the spreading of electron density over a larger volume of space. Electron shells are identified by the principal quantum numbers 1, 2, 3, and so forth. Each shell can contain up to 2n2 electrons, where n is the number of the shell. Thus, the first shell can contain 2 electrons; the second, 8 electrons; the third, 18 electrons; the fourth, 32 electrons; and so on (Table 1.1). Electrons in the first shell are nearest to the positively charged nucleus and are held most strongly by it; these electrons are lowest in energy. Electrons in higher-numbered shells are farther from the positively charged nucleus and are held less strongly. Shells are divided into subshells designated by the letters s, p, d, and f, and within these subshells, electrons are grouped in orbitals (Table 1.2). An orbital is a region of space that can hold two electrons and has a specific quantized energy. The first shell contains a single orbital called a 1s orbital. The second shell contains one s orbital and three p orbitals. The three 2p orbitals reflect orthogonal angular momentum states in three-dimensional space. Orthogonal in this context results in 90° angles between the orbitals, but in all cases, orthogonal also means that the orbitals have no net overlap. As a point of reference, to discuss the 2p orthogonal orbitals, we consider them to be directed along the x-, y-, and z-axes and give them designations, 2px, 2py, and Shell A region of space around a nucleus that can be occupied by electrons, corresponding to a principal quantum number. Quantized Having discrete values for energy and momentum. Delocalization The spreading of electron density over a larger volume of space. Orbital A region of space that can hold two electrons. Orthogonal Having no net overlap. Nucleus containing neutrons and protons Extranuclear space containing electrons 10–10 m Figure 1.1 A schematic view of an atom. Most of the mass of an atom is concentrated in its small, dense nucleus. Table 1.1 Distribution of Electrons in Shells Shell Number of Electrons Shell Can Hold Relative Energies of Electrons in These Shells 4 32 3 18 2 8 1 2 lower higher Copyright 2013 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
1.1 Table 1.2 Distribution of Orbitals in Shells of Atoms Orbitals Contained in That Shell 3s.plus five 3dorbitals 2x,2P29,2p 1s s are shown in Figures. A.Electron Configuration of Atoms elton configurlion fr occupy.Every atom has an ir finite numb er of pos mest en We determine the state electron round-state electron configuration of an atom by using the following three rules. configuration Rule 1:The Aufbau ("Build-Up)Principle.Orbitals fill in order of increasing energy from lowest to highest.In this course,we are concemed primarily with the elements Rule 2:The Pauli Exclusion Principle. The upy an o and that aheKchsioweeecpe eir spins must be paire h as the earth has a spin.electrons have a quantum mechanical pro perty referred to as spin.And just as the earth has magnetic north(N)and south(S)poles,so do elec- When filling orbitals with electrons,place no more than two in an orbital.For ex- ample,with four electrons,the Is and 2s orbitals are filled and are written s2.With tional si is written 2p2p s of equal energ ough elec toall of them Rule 3:Hund's Rule.Hund's rule has two parts.The first part states that when orbitals of equal energy(called degenerate)are available but there are not enough added to any Unless otherwise noted all art on this page Cengage Leaming 2013
Unless otherwise noted all art on this page © Cengage Learning 2013 3 1.1 Electronic Structure of Atoms 2pz . The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals. The shapes of s and p orbitals are shown in Figures 1.8 and 1.9 and are described in more detail in Section 1.6B. A. Electron Configuration of Atoms The electron configuration of an atom is a description of the orbitals its electrons occupy. Every atom has an infinite number of possible electron configurations. At this stage, we are concerned primarily with the ground-state electron configuration— the electron configuration of lowest energy. We determine the ground-state electron configuration of an atom by using the following three rules. Rule 1: The Aufbau (“Build-Up”) Principle. Orbitals fill in order of increasing energy, from lowest to highest. In this course, we are concerned primarily with the elements of the first, second, and third periods of the Periodic Table. Orbitals fill in the order 1s, 2s, 2p, 3s, 3p, and so on. Rule 2: The Pauli Exclusion Principle. The Pauli exclusion principle requires that only two electrons can occupy an orbital and that their spins must be paired. To understand what it means to have paired spins, recall from general chemistry that just as the earth has a spin, electrons have a quantum mechanical property referred to as spin. And just as the earth has magnetic north (N) and south (S) poles, so do electrons. As described by quantum mechanics, a given electron can exist in only two different spin states. Two electrons with opposite spins are said to have paired spins. When their tiny magnetic elds are aligned N-S, the electron spins are paired N S S N The quantum mechanical property of spin gives an electron a tiny magnetic eld When filling orbitals with electrons, place no more than two in an orbital. For example, with four electrons, the 1s and 2s orbitals are filled and are written 1s 2 2s 2. With an additional six electrons, the set of three 2p orbitals is filled and is written 2px 2 2py 2 2pz 2. Alternatively, a filled set of three 2p orbitals may be written 2p6. Rule 3: Hund’s Rule. Hund’s rule has two parts. The first part states that when orbitals of equal energy (called degenerate) are available but there are not enough Ground-state electron configuration The lowest-energy electron configuration for an atom or a molecule. Aufbau principle Orbitals fill in order of increasing energy, from lowest to highest. Pauli exclusion principle No more than two electrons may be present in an orbital. If two electrons are present, their spins must be paired. Hund’s rule When orbitals of equal energy are available but there are not enough electrons to fill all of them completely, one electron is put in each before a second electron is added to any. Table 1.2 Distribution of Orbitals in Shells Shell Orbitals Contained in That Shell 3 3s, 3px, 3py, 3pz , plus five 3d orbitals 2 2s, 2px, 2py, 2pz 1 1s Copyright 2013 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
Chapter Bonding and Shapes of Molecules Ground-State Electron Configurations for First Period" Second Period Third Period H 1 Li 3 [He]2s Na 11 [Ne]3s He 2 Is Be 4 [He]22 Mg 12 [Ne]3s B 5 [Hel 2s2p Al 13 INel 330 6 [He]2p 14 [Ne]33p 7[He2s22 15 [Ne]3s23p O 8 [Hel22 S 16 [Ne]3s3p 9 [He] Cl 17 [Ne]393p* Ne 10 [He]22 Ar 18 [Ne]3923p isted by symbolatomicnumberndsmfiedound-stateiurtio cec of them isadded to each orbital nd d ruestates that the spins of the ectron the aligned.Recall that electrons have a negative charge;partially filling orbitals as much repu s an bital tothe2Ponsasoaeeoaaeohe2poaCatooreanpe】 le,has atively,it may be ent 1 that pictorially des ration.For ple,the ngn level diagram for the electron configuration of carbon,,shows three en ach for the s,25, nd 2p orbital Moving up in the diagr am means the ls d the p exclusion principle tells us to pair the two electrons in each orbital (shown as arrows with opposing directions).The remaining two electrons are left to go into the 2p level,and b such orbita the second part of Hune rule tel us to place in throns in wn a 11 this chapter to explain bonding and throughout the book when discussing relative energies of orbitals. 2+ 1s
Unless otherwise noted all art on this page © Cengage Learning 2013 Chapter 1 Covalent Bonding and Shapes of Molecules 4 electrons to fill all of them completely, then one electron is added to each orbital before a second electron is added to any one of them. The second part of Hund’s rule states that the spins of the single electrons in the degenerate orbitals should be aligned. Recall that electrons have a negative charge; partially filling orbitals as much as possible minimizes electrostatic repulsion between electrons. After the 1s and 2s orbitals are filled with four electrons, a fifth electron is added to the 2px orbital, a sixth to the 2py orbital, and a seventh to the 2pz orbital. Only after each 2p orbital contains one electron is a second electron added to the 2px orbital. Carbon, for example, has six electrons, and its ground-state electron configuration is 1s2 2s2 2px 1 2py 1 2pz 0. Alternatively, it may be simplified to 1s2 2s2 2p2. Table 1.3 shows ground-state electron configurations of the first 18 elements of the Periodic Table. Chemists routinely write energy-level diagrams that pictorially designate where electrons are placed in an electron configuration. For example, the energylevel diagram for the electron configuration of carbon, 1s2, 2s2, 2p2, shows three energy levels, one each for the 1s, 2s, and 2p orbitals. Moving up in the diagram means higher energy. Electrons in these diagrams are drawn as arrows. The Aufbau principle tells us to place the first four electrons in the 1s and 2s orbitals, and the Pauli exclusion principle tells us to pair the two electrons in each orbital (shown as arrows with opposing directions). The remaining two electrons are left to go into the 2p level, and because there are three such orbitals, the second part of Hund’s rule tells us to place these electrons in different orbitals with their spins aligned (shown as arrows pointing in the same direction). We will use energy-level diagrams later in this chapter to explain bonding and throughout the book when discussing relative energies of orbitals. Energy 2s 1s 2p Energy level diagram for carbon Table 1.3 Ground-State Electron Configurations for Elements 1–18 First Period* Second Period Third Period H 1 1s1 Li 3 [He] 2s1 Na 11 [Ne] 3s1 He 2 1s2 Be 4 [He] 2s2 Mg 12 [Ne] 3s2 B 5 [He] 2s22p1 Al 13 [Ne] 3s23p1 C 6 [He] 2s22p2 Si 14 [Ne] 3s23p2 N 7 [He] 2s22p3 P 15 [Ne] 3s23p3 O 8 [He] 2s22p4 S 16 [Ne] 3s23p4 F 9 [He] 2s22p5 Cl 17 [Ne] 3s23p5 Ne 10 [He] 2s22p6 Ar 18 [Ne] 3s23p6 *Elements are listed by symbol, atomic number, and simplified ground-state electron configuration. Copyright 2013 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
1.1 Example 1.1 Electron Configurations of Atoms Write the ground-state electron configuration for each element showing the oc- cupancy of each p orbital.For(c),write the energy-level diagram. (a)Lithium (b)Oxygen (c)Chlorine Solution (a)Lithium (atomic number 3):1s22s (b)Oxygen (atomic number 8):p (c)Chlorine (atomic number 17:sppp33p23p23p p+北十 3 即+北什 2 1 Problem 1.1 Write and compare the ground-state electron configurations for each pair of elements. (a)Carbon and silicon (b)Oxygen and sulfur (c)Nitrogen and phosphorus B.The Concept of Energy hergy-level d the 2s and asking "How is energy defined? you may be Energy is the ability to do work.The higher in energy an entity is,the more work i erform.If you h an object above t do work 6 din the al ntialg the can h sed why Potential er stores and the greater the impact the object will have when it hits the ground. nhe2 stor tional a bject to ting sta e gro h th example.thousands of miles above the earth the obicct has incredibly large potential energy and could wreak serious damage to a building if dropped.But at that distance it isr asy to remove the object farther from the earth because the gravita- We can eralize this ovamnle to chemical structures Instable structures nos. sess enere Unless otherwise noted all art on this page Cengage Leaming 2013
Unless otherwise noted all art on this page © Cengage Learning 2013 5 1.1 Electronic Structure of Atoms Example 1.1 | Electron Configurations Write the ground-state electron configuration for each element showing the occupancy of each p orbital. For (c), write the energy-level diagram. (a) Lithium (b) Oxygen (c) Chlorine Solution (a) Lithium (atomic number 3): 1s2 2s1 (b) Oxygen (atomic number 8): 1s2 2s2 2px 2 2py 1 2pz 1 (c) Chlorine (atomic number 17): 1s2 2s22px 2 2py 2 2pz 23s2 3px 2 3py 2 3pz 1 1s 2s 2p 3s 3p Energy Energy level diagram for chlorine Problem 1.1 Write and compare the ground-state electron configurations for each pair of elements. (a) Carbon and silicon (b) Oxygen and sulfur (c) Nitrogen and phosphorus B. The Concept of Energy In the discussion of energy-level diagrams, the lines were drawn on the diagram to depict relative energy. In the energy-level diagram for carbon, the 1s level is the reference and the 2s and 2p levels are placed higher on the diagram relative to it. But you may be asking, “How is energy defined?” Energy is the ability to do work. The higher in energy an entity is, the more work it can perform. If you hold an object above the ground, it is unstable relative to when it is lying on the ground. You expend energy lifting the object, and this energy is stored in the object as potential energy. The potential energy can be released when the object is released. The higher you hold the object, the more energy the object stores and the greater the impact the object will have when it hits the ground. The force that restores the object to its resting state on the ground is the gravitational attraction of the object to the earth. Interestingly, the farther the object is from the earth, the easier it is to take the object even farther from the earth. As an extreme example, thousands of miles above the earth the object has incredibly large potential energy and could wreak serious damage to a building if dropped. But at that distance, it is relatively easy to remove the object farther from the earth because the gravitational attraction is weak. We can generalize this example to chemical structures. Unstable structures possess energy waiting to be released. When a structure is higher in energy, the more energy it has stored. When that energy is released, work can be done. In chemistry, Energy The ability to do work. Potential energy The energy that can be released if given an opportunity. Copyright 2013 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