Introduction 3 own bodies,are reactions of organic compounds.Most of the compounds found in nature-those that we rely on for food,clothing (cotton,wool,silk).and energy (natural gas,petroleum)are organic compounds as well synthetic rubber.and even things such as compact discs and Super Glue.And most importantly. almost all commonly prescribed drugs are synthetic organic compounds. Some synthetic organic compounds prevent shor of cotton and wool just to provide enough material to clothe us.Other synthetic organic compounds provide us with materials we would not have-Teflon,Plexiglas.Kevlar-if we had only naturally .there are about 6mllion known organic th Po The answer lies in carbon's position in the periodic table.Carbon is in the center of the second row of elements.We will see that the atoms helne e n cabon is in the middleit shareselectrons u hn cor the second row of the periodic table chemical properties simply by sharing electrons. compou Natural Versus Synthetic Organic Compounds -those made in the laboratory.Yet when a chemist synthesizes a compound,such as penicillin or morphine,the compound s the same in all respects as the compound synthesized in nature.Some an ev atotpencinaoeetptpocehe0epetesponshatasecaR fraction of the population experiences from naturally produced penicillin or that do not have the bacterial resistance of the naturally produced antibiotic nalogues of morphine that have the same pain-killing effects but,unlike morphine,are not habit-forming Most commercial morphine is obtained from opium,the juice extracted for es poppy shc e is the synthesis has an extremely pungent odor:dogs used by dru ment agencies are trained to recognize this odor (Section 15.16).Nearly afield of poppies inAfghanistan otheorls supply of eroin comsfrom the poppy elds
Introduction 3 own bodies, are reactions of organic compounds. Most of the compounds found in nature—those that we rely on for food, clothing (cotton, wool, silk), and energy (natural gas, petroleum)—are organic compounds as well. Organic compounds are not limited to those found in nature. Chemists have learned how to synthesize millions of organic compounds not found in nature, including synthetic fabrics, plastics, synthetic rubber, and even things such as compact discs and Super Glue. And most importantly, almost all commonly prescribed drugs are synthetic organic compounds. Some synthetic organic compounds prevent shortages of naturally occurring compounds. For example, it has been estimated that if synthetic materials—nylon, polyester, Lycra—were not available for clothing, all of the arable land in the United States would have to be used for the production of cotton and wool just to provide enough material to clothe us. Other synthetic organic compounds provide us with materials we would not have—Teflon, Plexiglas, Kevlar—if we had only naturally occurring organic compounds. Currently, there are about 16 million known organic compounds, and many more are possible that we cannot even imagine today. Why are there so many carbon-containing compounds? The answer lies in carbon’s position in the periodic table. Carbon is in the center of the second row of elements. We will see that the atoms to the left of carbon have a tendency to give up electrons, whereas the atoms to the right have a tendency to accept electrons (Section 1.3). Li Be B C the second row of the periodic table N O F carbon is in the middle—it shares electrons Because carbon is in the middle, it neither readily gives up nor readily accepts electrons. Instead, it shares electrons. Carbon can share electrons with several kinds of atoms as well as with other carbon atoms. Consequently, carbon forms millions of stable compounds with a wide range of chemical properties simply by sharing electrons. Natural Versus Synthetic Organic Compounds It is a popular belief that natural substances—those made in nature—are superior to synthetic ones—those made in the laboratory. Yet when a chemist synthesizes a compound, such as penicillin or morphine, the compound is the same in all respects as the compound synthesized in nature. Sometimes chemists can even improve on nature. For example, chemists have synthesized analogues of penicillin—compounds with structures similar to that of penicillin—that do not produce the allergic responses that a significant fraction of the population experiences from naturally produced penicillin or that do not have the bacterial resistance of the naturally produced antibiotic (Section 15.11). Chemists have also synthesized analogues of morphine that have the same pain-killing effects but, unlike morphine, are not habit-forming. Most commercial morphine is obtained from opium, the juice extracted from the species of poppy shown in the photo. Morphine is the starting material for the synthesis of heroin. One of the side products formed in the synthesis has an extremely pungent odor; dogs used by drug enforcement agencies are trained to recognize this odor (Section 15.16). Nearly three-quarters of the world’s supply of heroin comes from the poppy fields of Afghanistan. a field of poppies in Afghanistan
CHAPTER1 Remembering General Chemistry:Electronic Structure and Bonding organic compoun d aks depends on the that are shared.whic or che 1.1 THE STRUCTURE OF AN ATOM sists of a tiny dense nucle nded by electro s th The n relatively large volume of space around the nucleus called an electron cloud.The nucleus contains positively charged protons and uncharged neutrons,so it is positively charged.The electrons ged.The amount o Like anything that moves clectrns have kintic e ed atom m e number ust he the (nucleus(protons+neutrons) energy counteracts the attractive force of the positively charged protons that pull the negatively charged electrons toward the nucleus ondneoofcappooa ame ma and are about 1800 ti the is in its atom.however.is occupied by its electron cloud.This is where our focus will be because it is the electrons that form chemical bonds. The atomic number of an atom is the number of protons in its nucleus.The atomic number in an atom of a particular element never changes The mass number of an atom is the sum its protons and neutrons.Although all carbon atom have the no not all b ve tne sa electron cloud neutrons giving them a mass an atom number of 13.These two different kinds of carbon atoms(CandC)are called isotopes. otepiohaenhbe ooenumbers 12 13 14 isotopes of carbon Carbon also eutrons.This it takes for one-half of the nuclei to decay.)As long as a plant or an animal is alive.theC that is lost through chatioohrough radicactive deeay.Therefor the nving organism can )/12 of the mass ofC.the mass ofCis 12.0000amu:the massof is 13.0035 amu.Therefore,the atomic mass of carbon is 12.011 amu because(.9889 12.0000) (0.0111 ×13.0035) 12.011.The molecular mass is the sum of the atomic masses of all the toms held toge atoms in the molecule
4 CHAPTER 1 Remembering General Chemistry: Electronic Structure and Bonding When we study organic chemistry, we learn how organic compounds react. Organic compounds consist of atoms held together by covalent bonds. When an organic compound reacts, some of these covalent bonds break and some new covalent bonds form. Covalent bonds form when two atoms share electrons, and they break when two atoms no longer share electrons. How easily a covalent bond forms or breaks depends on the electrons that are shared, which, in turn, depends on the atoms to which the electrons belong. So if we are going to start our study of organic chemistry at the beginning, we must start with an understanding of the structure of an atom—what electrons an atom has and where they are located. 1.1 THE STRUCTURE OF AN ATOM An atom consists of a tiny dense nucleus surrounded by electrons that are spread throughout a relatively large volume of space around the nucleus called an electron cloud. The nucleus contains positively charged protons and uncharged neutrons, so it is positively charged. The electrons are negatively charged. The amount of positive charge on a proton equals the amount of negative charge on an electron. Therefore, the number of protons and the number of electrons in an uncharged atom must be the same. Electrons move continuously. Like anything that moves, electrons have kinetic energy, and this energy counteracts the attractive force of the positively charged protons that pull the negatively charged electrons toward the nucleus. Protons and neutrons have approximately the same mass and are about 1800 times more massive than an electron. Most of the mass of an atom, therefore, is in its nucleus. Most of the volume of an atom, however, is occupied by its electron cloud. This is where our focus will be because it is the electrons that form chemical bonds. The atomic number of an atom is the number of protons in its nucleus. The atomic number is unique to a particular element. For example, the atomic number of carbon is 6, which means that all uncharged carbon atoms have six protons and six electrons. Although atoms can gain electrons and become negatively charged or lose electrons and become positively charged, the number of protons in an atom of a particular element never changes. The mass number of an atom is the sum of its protons and neutrons. Although all carbon atoms have the same atomic number, they do not all have the same mass number. Why? Because carbon atoms can have varying numbers of neutrons. For example, 98.89% of all carbon atoms have six neutrons—giving them a mass number of 12—and 1.11% have seven neutrons—giving them a mass number of 13. These two different kinds of carbon atoms (12C and 13C) are called isotopes. 13C 6 12C 6 14C 6 isotopes of carbon isotopes have the same atomic number isotopes have different mass numbers Carbon also contains a trace amount of 14C, which has six protons and eight neutrons. This isotope of carbon is radioactive, decaying with a half-life of 5730 years. (The half-life is the time it takes for one-half of the nuclei to decay.) As long as a plant or an animal is alive, the 14C that is lost through exhalation or excretion is constantly replenished. When it dies, however, it no longer ingests 14C. Consequently, its 14C is slowly lost through radioactive decay. Therefore, the age of a substance derived from a living organism can be determined by its 14C content. The atomic mass is the weighted average of the isotopes in the element. Because an atomic mass unit (amu) is defined as exactly 1/12 of the mass of 12C, the mass of 12C is 12.0000 amu; the mass of 13C is 13.0035 amu. Therefore, the atomic mass of carbon is 12.011 amu because (0.9889 * 12.0000) + (0.0111 * 13.0035) = 12.011. The molecular mass is the sum of the atomic masses of all the atoms in the molecule. The nucleus contains positively charged protons and uncharged neutrons. The electrons are negatively charged. nucleus (protons + neutrons) an atom electron cloud atomic number = the number of protons in the nucleus mass number = the number of protons + the number of neutrons atomic mass = the weighted average mass of the isotopes in the element A molecule is a group of two or more atoms held together by bonds. molecular mass = the sum of the atomic masses of all the atoms in the molecule
1.2 How The Electrons in an Atom are Distributed 5 PROBLEM1◆ PROBLEM2◆ 2 Ar 3.c PROBLEM3◆ s.35Cl and Cl:75.77%of chlo is 35CL and 24 23%is Cl The Albert Einstein 2HOW THE ELECTRONS IN AN ATOM The bronze n Wa ARE DISTRIBUTED on,D.C measures 21feet from the For a long time.elec e a Fre ree ost im tant contributions to scie also have wave-like n relating mass and energy with a formula developed by Max Planck relating frequency and energy. and matte mechanics Quantum mechanics uses the same mathematical equations that describe the wave motion of a guitar string to characterize the motion of an electron around a nucleus.The version of quantum mechanics most useful to chemists was proposed by Erwin Schrodinger in 1926. Table 1.1 Electrons in the First Four Shells First shell Second shell Third shell Fourth shel Atomic orbitals p s.p.d s乃，d.f Number of atomic orbitals 1,3 135 13.5.7 Maximum number 8 18 32 of electrons .The first shell is the one closest to the nucleus.The second shell lies farther from the nucleus. The third and higher numbered shells lie even farther out. .The shells contain subshells known as atomic orbitals.We will see that an atomic orbital has a characteristic shape and energy and occupies a characteristic volume of space(Section 1.5). atomic orbitals D orbitals are orbitals .A maximum of two electrons can coexist in an atomic orbital.(See the Pauli exclusion principle on p.6.)Th .9,and 16 atomic orbitals
1.2 How The Electrons in an Atom are Distributed 5 Albert Einstein The bronze sculpture of Albert Einstein, on the grounds of the National Academy of Sciences in Washington, D.C., measures 21 feet from the top of the head to the tip of the feet and weighs 7000 pounds. In his left hand, Einstein holds the mathematical equations that represent his three most important contributions to science: the photoelectric effect, the equivalency of energy and matter, and the theory of relativity. At his feet is a map of the sky. Degenerate orbitals are orbitals that have the same energy. PROBLEM 1 ♦ Oxygen has three isotopes, 16O, 17O, and 18O. The atomic number of oxygen is 8. How many protons and neutrons does each of the isotopes have? PROBLEM 2 ♦ a. How many protons do the following species have? (See the periodic table inside the back cover of this book.) b. How many electrons does each have? 1. Na+ 2. Ar 3. ClPROBLEM 3 ♦ Chlorine has two isotopes, 35Cl and 37Cl; 75.77% of chlorine is 35Cl, and 24.23% is 37Cl. The atomic mass of 35Cl is 34.969 amu, and the atomic mass of 37Cl is 36.966 amu. What is the atomic weight of chlorine? 1.2 HOW THE ELECTRONS IN AN ATOM ARE DISTRIBUTED For a long time, electrons were perceived to be particles—infinitesimal “planets” that orbit the nucleus of an atom. In 1924, however, Louis de Broglie, a French physicist, showed that electrons also have wave-like properties. He did this by combining a formula developed by Albert Einstein relating mass and energy with a formula developed by Max Planck relating frequency and energy. The realization that electrons have wave-like properties spurred physicists to propose a mathematical concept known as quantum mechanics to describe the motion of an electron around a nucleus. Quantum mechanics uses the same mathematical equations that describe the wave motion of a guitar string to characterize the motion of an electron around a nucleus. The version of quantum mechanics most useful to chemists was proposed by Erwin Schrödinger in 1926. According to Schrödinger, the electrons in an atom can be thought of as occupying a set of concentric shells that surround the nucleus (Table 1.1). Table 1.1 Distribution of Electrons in the First Four Shells First shell Second shell Third shell Fourth shell Atomic orbitals s s, p s, p, d s, p, d, f Number of atomic orbitals 1 1, 3 1, 3, 5 1, 3, 5, 7 Maximum number of electrons 2 8 18 32 ■ The first shell is the one closest to the nucleus. The second shell lies farther from the nucleus. The third and higher numbered shells lie even farther out. ■ The shells contain subshells known as atomic orbitals. We will see that an atomic orbital has a characteristic shape and energy and occupies a characteristic volume of space (Section 1.5). ■ Each shell contains one s atomic orbital. Each second and higher shell—in addition to its s atomic orbital—contains three degenerate p atomic orbitals. Degenerate orbitals are orbitals that have the same energy. The third and higher shells—in addition to their s and p atomic orbitals—contain five degenerate d atomic orbitals, and the fourth and higher shells also contain seven degenerate f atomic orbitals. ■ A maximum of two electrons can coexist in an atomic orbital. (See the Pauli exclusion principle on p. 6.) Therefore, the first four shells, with 1, 4, 9, and 16 atomic orbitals, respectively, can contain a maximum of 2, 8, 18, and 32 electrons. In our study of organic chemistry, we will be concerned primarily with atoms that have electrons only in the first two shells
6 CHAPTER1 Remembering General Chemistry:Electronic Structure and Bonding Ground-State Electronic Configuration m in the ppid ed ste d vdt ect The ground-state electronic configurations of the smallest atoms are shown in Table 1.2.(Each arrow- -whether pointing up or down-represents one electron. Table 1.2 The the Smallest Atoms Atom Name of element Atomic number Is 2p.2p,2p: H Hydrogen 2 Lithium 3 Be Beryllium 4 ↑ R Boron ↑ ↑ Carbon Nitrogen 0 Oxygen 8 ↑ ↑ Fluorin 0 ↑ Ne 10 Na Sodium 11 Three rules specify which atomic orbitals an atom's electrons occupy: 1.The aufbau principle (auau is German for"building up)states that an electron always goes into the available orbital with the lowest energy. r the atomicorbit is to the 2s orbital.whichs d closer to th When comparing atomic orbitals in the same shell.we see that an s orbital is lower in energy than ap orbital,and ap orbital is lower in energy than a d orbital. relative energies of atomic orbital lowest<2s 2p<3s <3p <3d-highest energy 2.The Pauli exclusion principle states that ne,Terynocrnachaomtcorbialamnd This is called an exclu wsto sin electrons to atomic oritals for atoms that contan one two,t .The single electron of a hydrogen atom occupies a Is orbital The second electron of a helium atom fills the Is orbital. The third electron of a lithium atom occupies a 2s orbital The fourth electron of a beryllium atom fills Fdissire tr)Bcause Before we can discuss atoms containing six or more electrons,we need the third rule
6 CHAPTER 1 Remembering General Chemistry: Electronic Structure and Bonding Ground-State Electronic Configuration The ground-state electronic configuration of an atom describes the atomic orbitals occupied by the atom’s electrons when they are all in the available orbitals with the lowest energy. If energy is applied to an atom in the ground state, one or more electrons can jump into a higher-energy orbital. The atom then would be in an excited state and have an excited-state electronic configuration. The ground-state electronic configurations of the smallest atoms are shown in Table 1.2. (Each arrow—whether pointing up or down—represents one electron.) Atom Name of element Atomic number 1s 2s 2px 2py 2pz 3s H Hydrogen 1 c He Helium 2 cT Li Lithium 3 cT c Be Beryllium 4 cT cT B Boron 5 cT cT c C Carbon 6 cT cT c c N Nitrogen 7 cT cT c c c O Oxygen 8 cT cT cT c c F Fluorine 9 cT cT cT cT c Ne Neon 10 cT cT cT cT cT Na Sodium 11 cT cT cT cT cT c Three rules specify which atomic orbitals an atom’s electrons occupy: 1. The aufbau principle (aufbau is German for “building up”) states that an electron always goes into the available orbital with the lowest energy. When using the aufbau principle rule, it is important to remember that the closer the atomic orbital is to the nucleus, the lower is its energy. Because the 1s orbital is closer to the nucleus, it is lower in energy than the 2s orbital, which is lower in energy—and closer to the nucleus—than the 3s orbital. When comparing atomic orbitals in the same shell, we see that an s orbital is lower in energy than a p orbital, and a p orbital is lower in energy than a d orbital. relative energies of atomic orbitals lowest energy 1s < 2s < 2p < 3s < 3p < 3d highest energy 2. The Pauli exclusion principle states that no more than two electrons can occupy each atomic orbital, and the two electrons must be of opposite spin. This is called an exclusion principle because it limits the number of electrons that can occupy an atomic orbital and, therefore, any particular shell. (Notice in Table 1.2 that opposite spins are designated by c and T.) These first two rules allow us to assign electrons to atomic orbitals for atoms that contain one, two, three, four, or five electrons. ■ The single electron of a hydrogen atom occupies a 1s orbital. ■ The second electron of a helium atom fills the 1s orbital. ■ The third electron of a lithium atom occupies a 2s orbital. ■ The fourth electron of a beryllium atom fills the 2s orbital. ■ The fifth electron of a boron atom occupies one of the 2p orbitals. (The subscripts x, y, and z distinguish the three 2p orbitals.) Because the three p orbitals are degenerate, the electron can be put into any one of them. Before we can discuss atoms containing six or more electrons, we need the third rule. Table 1.2 The Electronic Configurations of the Smallest Atoms
1.3 Covalent Bonds 7 3.Hund's rule states that In this way.electron repulsion is minimized. Therefore.the sixth electron of a carbon atom goes into an empty 2p orbital,rather than pairing up with the electron already occupying a 2p orbital (see Table 1.2). .There is one more empty 2p orbital,so that is where nitrogen's seventh electron goes. .The eighth electron of an oxygen atom pairs up with an electron occupying a 2p orbital rather than going into the higher-energy 3s orbital. The locations of the electrons in the remaining elements can be assigned using these three rules. Valence and Core Electrons The maior factor that determines the chemical behay it has.Valence electrons are electrons in an atom's outermost shell Electrons in inner shells (below the outermost shell)are called core electrons.For example,carbon has four valence electrons and two core electrons (Tabl 1.2).Valence electr s partic mn of ntenchercalbondtnegrcoeeectosdbnot se they ,are in The chemical behavior of the same column because each has one valence electron. PROBLEM4◆ How many valence electrons do the following atoms have? a.boron b.nitrogen c.oxygen d.且orine PROBLEM5◆ Wite theo for chorn (aomic number bromine (atomic romine,and iodine have? PROBLEM 6 Look a.carbon and silicor c.nitrogen and phosphorus b.oxygen and sulfur d.magnesium and calcium 1.3 COVALENT BONDS Now that you know about the electronic configuration of atoms.let's now look at why atoms come together to form bonds.In explaining why atoms form bonds,G.N.Lewis proposed that hell th (even though hydrogen needs only two electrons to achieve a filled outer shell). Achieving a Filled Outer Shell by Losing or Gaining Electrons Sodium (Na)has a single electron in its3s orbital:so it.too.loses an electron easily
1.3 Covalent Bonds 7 Valence electrons are electrons in the outermost shell. Core electrons are electrons in inner shells. The chemical behavior of an element depends on its electronic configuration. 3. Hund’s rule states that when there are two or more atomic orbitals with the same energy, an electron will occupy an empty orbital before it will pair up with another electron. In this way, electron repulsion is minimized. ■ Therefore, the sixth electron of a carbon atom goes into an empty 2p orbital, rather than pairing up with the electron already occupying a 2p orbital (see Table 1.2). ■ There is one more empty 2p orbital, so that is where nitrogen’s seventh electron goes. ■ The eighth electron of an oxygen atom pairs up with an electron occupying a 2p orbital rather than going into the higher-energy 3s orbital. The locations of the electrons in the remaining elements can be assigned using these three rules. Valence and Core Electrons The major factor that determines the chemical behavior of an element is the number of valence electrons it has. Valence electrons are electrons in an atom’s outermost shell. Electrons in inner shells (below the outermost shell) are called core electrons. For example, carbon has four valence electrons and two core electrons (Table 1.2). Valence electrons participate in chemical bonding; core electrons do not. Elements in the same column of the periodic table have similar chemical properties because they have the same number of valence electrons. If you examine the periodic table inside the back cover of this book, you will see that lithium and sodium, which have similar chemical properties, are in the same column because each has one valence electron. PROBLEM 4 ♦ How many valence electrons do the following atoms have? a. boron b. nitrogen c. oxygen d. fluorine PROBLEM 5 ♦ a. Write the ground-state electronic configuration for chlorine (atomic number 17), bromine (atomic number 35), and iodine (atomic number 53). b. How many valence electrons do chlorine, bromine, and iodine have? PROBLEM 6 Look at the relative positions of each pair of atoms listed here in the periodic table. How many core electrons does each have? How many valence electrons does each have? a. carbon and silicon b. oxygen and sulfur c. nitrogen and phosphorus d. magnesium and calcium 1.3 COVALENT BONDS Now that you know about the electronic configuration of atoms, let’s now look at why atoms come together to form bonds. In explaining why atoms form bonds, G. N. Lewis proposed that an atom is most stable if its outer shell is either filled or contains eight electrons, and it has no electrons of higher energy. According to Lewis’s theory, an atom will give up, accept, or share electrons to achieve a filled outer shell or an outer shell that contains eight electrons. This theory has come to be called the octet rule (even though hydrogen needs only two electrons to achieve a filled outer shell). Achieving a Filled Outer Shell by Losing or Gaining Electrons Lithium (Li) has a single electron in its 2s orbital. If it loses this electron, lithium ends up with a filled outer shell—a stable configuration. Lithium, therefore, loses an electron relatively easily. Sodium (Na) has a single electron in its 3s orbital; so it, too, loses an electron easily