BRUI01-001_059r4 20-03-2003 2:58 PM Page 1 The first twochaptersof the text cover a An Introduction P variety of topics that you need to get started A with your study of organic chemistry. to the Study R Chapter 1 reviews the topics from general chemistry of Organic that will be important to your study of organic chemistry. The chapter starts with a description of the structure of Chemistry T atoms and then proceeds to a description of the structure of molecules. Molecular orbital theory is introduced. Acid–base chemistry, which is central to understanding O many organic reactions,is reviewed. You will see how the structure of a molecule affects its acidity and how the acidity of a solution affects molecular structure. N To discuss organic compounds,you must be able to name them and visualize their structures when you read or hear E their names. In Chapter 2, you will learn how to name five different classes of organic compounds. This will give you a good understanding of the basic rules followed Chapter 1 in naming compounds. Because the compounds exam- ined in the chapter are either the reactants or the products Electronic Structure and Bonding of many of the reactions presented in the next 10 chap- • Acids and Bases ters, you will have the opportunity to review the nomen- clature of these compounds as you proceed through those Chapter 2 chapters. The structures and physical properties of these An Introduction to Organic compounds will be compared and contrasted, which Compounds:Nomenclature, makes learning about them a little easier than if each Physical Properties, and compound were presented separately. Because organic chemistry is a study of compounds that contain carbon, Representation of Structure the last part of Chapter 2 discusses the spatial arrange- ment of the atoms in both chains and rings of carbon atoms. 1 BRUI01-001_059r4 20-03-2003 2:58 PM Page 2 1 Electronic Structure and Bonding • Acids and Bases Ethane Ethene T o stay alive, early humans must have been able to tell the difference between twokinds of materials in their world. “You can live on roots and berries,” they might have Jöns Jakob Berzelius (1779–1848) said, “but you can’t live on dirt. You can not only coined the terms “organic” stay warm by burning tree branches, but and “inorganic,”but also invented you can’t burn rocks.” Ethyne the system of chemical symbols still By the eighteenth century,scientists thought they used today. He published the first list had grasped the nature of that difference, and in 1807, Jöns Jakob Berzelius gave of accurate atomic weights and names to the two kinds of materials. Compounds derived from living organisms were proposed the idea that atoms carry believed to contain an unmeasurable vital force—the essence of life. These he called an electric charge. He purified or discovered the elements cerium, “organic.” Compounds derived from minerals—those lacking that vital force—were selenium,silicon,thorium,titanium, “inorganic.” and zirconium. Because chemists could not create life in the laboratory,they assumed they could not create compounds with a vital force. With this mind-set,you can imagine how surprised chemists were in 1828 when Friedrich Wöhler produced urea—a compound known to German chemist Friedrich Wöhler be excreted by mammals—by heating ammonium cyanate,an inorganic mineral. (1800–1882)began his professional life as a physician and later became a professor of chemistry at the Uni- O versity of Göttingen. Wöhler codis- + − heat NH OCN C covered the fact that two different 4 chemicals could have the same mo- ammonium cyanate H2N NH2 urea lecular formula. He also developed methods of purifying aluminum—at the time,the most expensive metal on For the first time, an “organic” compound had been obtained from something other Earth—and beryllium. than a living organism and certainly without the aid of any kind of vital force. Clearly, chemists needed a new definition for “organic compounds.”Organic compoundsare now defined as compounds that contain carbon. Why is an entire branch of chemistry devoted to the study of carbon-containing compounds? We study organic chemistry because just about all of the molecules that 2 BRUI01-001_059r4 20-03-2003 2:58 PM Page 3 Section 1.1 The Structure of an Atom 3 make life possible—proteins, enzymes, vitamins, lipids, carbohydrates, and nucleic acids—contain carbon,so the chemical reactions that take place in living systems,in- cluding our own bodies, are organic reactions. Most of the compounds found in nature—those we rely on for food,medicine,clothing (cotton,wool,silk),and energy (natural gas, petroleum)—are organic as well. Important organic compounds are not, however, limited to the ones we find in nature. Chemists have learned to synthesize millions of organic compounds never found in nature, including synthetic fabrics, plastics, synthetic rubber, medicines, and even things like photographic film and Superglue. Many of these synthetic compounds prevent shortages of naturally occur- ring products. For example, it has been estimated that if synthetic materials 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. Currently, there are about 16 million known organic compounds, and many more are possible. What makes carbon so special? Why are there so many carbon-containing com- pounds? The answer lies in carbon’s position in the periodic table. Carbon is in the center of the second row of elements. 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 (Section1.3). Li Be B C N O F the second row of the periodic table Because carbon is in the middle, it neither readily gives up nor readily accepts elec- trons. Instead, it shares electrons. Carbon can share electrons with several different kinds of atoms,and it can also share electrons with other carbon atoms. Consequently, carbon is able to form millions of stable compounds with a wide range of chemical properties simply by sharing electrons. When we study organic chemistry, we study how organic compounds react. When an organic compound reacts,some old bonds break and some new bonds form. Bonds form when two atoms share electrons, and bonds break when two atoms no longer share electrons. How readily a bond forms and how easily it breaks depend on the par- ticular electrons that are shared,which,in turn,depend on the atoms to which the elec- trons 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. The nucleus con- tains positively charged protons and neutral neutrons, so it is positively charged. The electrons are negatively charged. Because the amount of positive charge on a proton equals the amount of negative charge on an electron,a neutral atom has an equal num- ber of protons and electrons. Atoms can gain electrons and thereby become negatively charged,or they can lose electrons and become positively charged. However,the num- ber of protons in an atom does not change. Protons and neutrons have approximately the same mass and are about 1800 times more massive than an electron. This means that most of the mass of an atom is in its nucleus. However,most of the volumeof an atom is occupied by its electrons,and that is where our focus will be because it is the electrons that form chemical bonds. BRUI01-001_059r4 20-03-2003 2:58 PM Page 4 4 CHAPTER 1 Electronic Structure and Bonding • Acids and Bases Louis Victor Pierre Raymond duc The atomic number of an atom equals the number of protons in its nucleus. The de Broglie (1892–1987)was born in atomic number is also the number of electrons that surround the nucleus of a neutral France and studied history at the atom. For example,the atomic number of carbon is 6,which means that a neutral car- Sorbonne. During World War I,he bon atom has six protons and six electrons. Because the number of protons in an atom was stationed in the Eiffel Tower as a does not change, the atomic number of a particular element is always the same—all radio engineer. Intrigued by his expo- carbon atoms have an atomic number of 6. sure to radio communications,he re- The mass numberof an atom is the sumof its protons and neutrons. Not all carbon turned to school after the war,earned atoms have the same mass number,because,even though they all have the same num- a Ph.D. in physics,and became a ber of protons, they do not all have the same number of neutrons. For example, professor of theoretical physics at the 98.89% of naturally occurring carbon atoms have six neutrons—giving them a mass Faculté des Sciences at the Sorbonne. He received the Nobel Prize in number of 12—and 1.11% have seven neutrons—giving them a mass number of 13. physics in 1929,five years after ob- These two different kinds of carbon atoms (12Cand 13C)are called isotopes. Isotopes taining his degree,for his work that have the same atomic number (i.e., the same number of protons), but different mass showed electrons to have properties numbers because they have different numbers of neutrons. The chemical properties of of both particles and waves. In 1945, isotopes of a given element are nearly identical. he became an adviser to the French Naturally occurring carbon also contains a trace amount of 14C,which has six pro- Atomic Energy Commissariat. tons 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 animal is alive, it takes in as much 14C as it excretes or exhales. When it dies, it no longer takes in 14C, so the 14C in the organism slowly decreases. Therefore,the age of an organic substance can be determined by its 14Ccontent. The atomic weight of a naturally occurring element is the average weighted mass of its atoms. Because an atomic mass unit (amu) is defined as exactly 1 12 of the mass of 12C, the atomic mass of 12C is 12.0000 amu; the atomic > mass of 13Cis 13.0034 amu. Therefore,the atomic weight of carbon is 12.011 amu 0.9889 * 12.0000 + 0.0111 * 13.0034 = 12.011 . The molecular weight is the 1 2 sum of the atomic weights of all the atoms in the molecule. PROBLEM 1(cid:2) Oxygen has three isotopes with mass numbers of 16, 17, and 18. The atomic number of oxygen is eight. How many protons and neutrons does each of the isotopes have? 1.2 The Distribution of Electrons in an Atom Electrons are moving continuously. Like anything that moves, electrons have kinetic energy, and this energy is what counters the attractive force of the positively charged protons that would otherwise pull the negatively charged electrons into the nucleus. Erwin Schrödinger (1887–1961) For a long time, electrons were perceived to be particles—infinitesimal “planets” or- was teaching physics at the Universi- biting the nucleus of an atom. In 1924, however, a French physicist named Louis de ty of Berlin when Hitler rose to Broglie showed that electrons also have wavelike properties. He did this by combining power. Although not Jewish, a formula developed by Einstein that relates mass and energy with a formula devel- Schrödinger left Germany to return oped by Planck relating frequency and energy. The realization that electrons have to his native Austria,only to see it wavelike properties spurred physicists to propose a mathematical concept known as taken over later by the Nazis. He quantum mechanics. moved to the School for Advanced Quantum mechanicsuses the same mathematical equations that describe the wave Studies in Dublin and then to Oxford motion of a guitar string to characterize the motion of an electron around a nucleus. University. In 1933,he shared the The version of quantum mechanics most useful to chemists was proposed by Erwin Nobel Prize in physics with Paul Schrödinger in 1926. According to Schrödinger, the behavior of each electron in an Dirac,a professor of physics at Cam- bridge University,for mathematical atom or a molecule can be described by a wave equation. The solutions to the work on quantum mechanics. Schrödingerequation are called wave functionsor orbitals. They tell us the energyof the electron and the volume of space around the nucleus where an electron is most An orbital tells us the energy of the likely to befound. electron and the volume of space According to quantum mechanics,the electrons in an atom can be thought of as oc- around the nucleus where an electron is most likely to be found. cupying a set of concentric shells that surround the nucleus. The first shell is the one BRUI01-001_059r4 20-03-2003 2:58 PM Page 5 Section 1.2 The Distribution of Electrons in an Atom 5 ALBERT EINSTEIN Albert Einstein (1879–1955) was born in Germany. When he was in high school, his father’s business failed and his family moved to Milan, Italy. Einstein had to stay behind because German law required compulsory military service after finishing high school. Einstein wanted to join his family in Italy. His high school mathematics teacher wrote a letter saying that Einstein could have a nervous breakdown without his family and also that there was nothing left to teach him. Eventually,Einstein was asked to leave the school because of his disruptive behavior. Popular folklore says he left because of poor grades in Latin and Greek,but his grades in those subjects were fine. Einstein was visiting the United States when Hitler came to power,so he accepted a position at the Institute for Advanced Study in Princeton, becoming a U.S. citizen in 1940. Although a lifelong pacifist,he wrote a letter to President Roosevelt warning of ominous advances in Ger- man nuclear research. This led to the creation of the Manhattan Project, which developed the atomic bomb and tested it in New Mexico in 1945. closest to the nucleus. The second shell lies farther from the nucleus,and even farther out lie the third and higher numbered shells. Each shell contains subshells known as atomic orbitals. Each atomic orbital has a characteristic shape and energy and occu- pies a characteristic volume of space,which is predicted by the Schrödinger equation. An important point to remember is that the closer the atomic orbital is to the nucleus, The closer the orbital is to the nucleus, the lower is its energy. the lower is its energy. The first shell consists of only an satomic orbital; the second shell consists of sand patomic orbitals; the third shell consists of s,p,and datomic orbitals; and the fourth and higher shells consist of s,p,d,and fatomic orbitals(Table1.1). Each shell contains one satomic orbital. The second and higher shells—in addition to their s orbital—each contain three degenerate p atomic orbitals. Degenerate orbitals are orbitals that have the same energy. The third and higher shells—in Table1.1 Distribution of Electrons in the First Four Shells That Surround the Nucleus 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 MAX KARL ERNST LUDWIG PLANCK Max Planck (1858–1947) was born in Germany,the son of a professor of civil law. He was a professor at the Universities of Munich (1880–1889) and Berlin (1889–1926). Two of his daughters died in childbirth,and one of his sons was killed in action in World War I. In 1918,Planck received the Nobel Prize in physics for his development of quantum theory. He be- came president of the Kaiser Wilhelm Society of Berlin—later renamed the Max Planck Society— in 1930. Planck felt that it was his duty to remain in Germany during the Nazi era, but he never supported the Nazi regime. He unsuccessfully interceded with Hitler on behalf of his Jewish col- leagues and,as a consequence,was forced to resign from the presidency of the Kaiser Wilhelm So- ciety in 1937. A second son was accused of taking part in the plot to kill Hitler and was executed. Planck lost his home to Allied bombings. He was rescued by Allied forces during the final days of the war. BRUI01-001_059r4 20-03-2003 2:58 PM Page 6 6 CHAPTER 1 Electronic Structure and Bonding • Acids and Bases addition to their s and p orbitals—also contain five degenerate d atomic orbitals, and the fourth and higher shells also contain seven degenerate fatomic orbitals. Because a maximum of two electrons can coexist in an atomic orbital (see the Pauli exclusion principle, below), the first shell, with only one atomic orbital, can contain no more than two electrons. The second shell, with four atomic orbitals—one s and three p— can have a total of eight electrons. Eighteen electrons can occupy the nine atomic orbitals—one s,three p,and five d—of the third shell,and 32 electrons can occupy the 16 atomic orbitals of the fourth shell. In studying organic chemistry, we will be con- cerned primarily with atoms that have electrons only in the first and second shells. The ground-state electronic configuration of an atom describes the orbitals occu- pied by the atom’s electrons when they are all in the available orbitals with the lowest en- ergy. 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 electronic configuration. The ground-state electronic configurations of the 11 smallest atoms are shown inTable1.2. (Each arrow—whether pointing up or down—represents one elec- tron.) The following principles are used to determine which orbitals electrons occupy: 1. The aufbau principle (aufbau is German for “building up”) tells us the first thing we need to know to be able to assign electrons to the various atomic or- bitals. According to this principle,an electron always goes into the available or- bital with the lowest energy. The relative energies of the atomic orbitals are as follows: 1s 6 2s 6 2p 6 3s 6 3p 6 4s 6 3d 6 4p 6 5s 6 4d 6 5p 6 6s 6 4f 6 5d 6 6p 6 7s 6 5f Because a 1satomic orbital is closer to the nucleus,it is lower in energy than a 2satomic orbital,which is lower in energy—and is closer to the nucleus—than a 3satomic orbital. Comparing atomic orbitals in the same shell,we see that an s atomic orbital is lower in energy than a patomic orbital,and a patomic orbital is lower in energy than a datomic orbital. 2. The Pauli exclusion principlestates that (a)no more than two electrons can oc- cupy each atomic orbital,and (b)the two electrons must be of opposite spin. It is called an exclusion principle because it states that only so many electrons can occupy any particular shell. Notice in Table1.2that spin in one direction is des- ignated by an upward-pointing arrow, and spin in the opposite direction by a downward-pointingarrow. TABLE 1.2 The Ground-State Electronic Configurations of the Smallest Atoms Name of Atomic Atom element number 1s 2s 2p 2p 2p 3s x y z H Hydrogen 1 He Helium 2 Li Lithium 3 Be Beryllium 4 As a teenager,Austrian Wolfgang B Boron 5 Pauli (1900–1958)wrote articles on C Carbon 6 relativity that caught the attention of Albert Einstein. Pauli went on to N Nitrogen 7 teach physics at the University of O Oxygen 8 Hamburg and at the Zurich Institute F Fluorine 9 of Technology. When World War II Ne Neon 10 broke out,he immigrated to the Unit- ed States,where he joined the Insti- Na Sodium 11 tute for Advanced Study at Princeton. BRUI01-001_059r4 20-03-2003 2:58 PM Page 7 Section 1.3 Ionic, Covalent, and Polar Bonds 7 From these first two rules,we can assign electrons to atomic orbitals for atoms that contain one,two,three,four,or five electrons. The single electron of a hydrogen atom Tutorial: occupies a 1satomic orbital,the second electron of a helium atom fills the 1satomic Electrons in orbitals orbital, the third electron of a lithium atom occupies a 2s atomic orbital, the fourth electron of a beryllium atom fills the 2satomic orbital,and the fifth electron of a boron Friedrich Hermann Hund atom occupies one of the 2patomic orbitals. (The subscripts x,y,and zdistinguish the (1896–1997)was born in Germany. three 2patomic orbitals.) Because the three porbitals are degenerate,the electron can He was a professor of physics at sev- be put into any one of them. Before we can continue to larger atoms—those contain- eral German universities,the last being the University of Göttingen. He ing six or more electrons—we need Hund’s rule: spent a year as a visiting professor at 3. Hund’s rulestates that when there are degenerate orbitals—two or more orbitals Harvard University. In February with the same energy—an electron will occupy an empty orbital before it will 1996,the University of Göttingen pair up with another electron. In this way,electron repulsion is minimized. The held a symposium to honor Hund on sixth electron of a carbon atom,therefore,goes into an empty 2patomic orbital, his 100th birthday. rather than pairing up with the electron already occupying a 2p atomic orbital. (See Table1.2.) The seventh electron of a nitrogen atom goes into an empty 2p atomic orbital,and the eighth electron of an oxygen atom pairs up with an elec- tron occupying a 2p atomic orbital rather than going into a higher energy 3s orbital. Using these three rules,the locations of the electrons in the remaining elements can be assigned. PROBLEM 2(cid:2) Potassium has an atomic number of 19 and one unpaired electron. What orbital does the unpaired electron occupy? PROBLEM 3(cid:2) Write electronic configurations for chlorine (atomic number 17),bromine (atomic number 35),and iodine (atomic number 53). 1.3 Ionic, Covalent, and Polar Bonds In trying to explain 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 in order to achieve a filled outer shell or an outer shell that contains eight electrons. This theory has come to be called the octet rule. Lithium (Li) has a single electron in its 2satomic orbital. If it loses this electron,the lithium atom ends up with a filled outer shell—a stable configuration. Removing an electron from an atom takes energy—called the ionization energy. Lithium has a rel- atively low ionization energy—the drive to achieve a filled outer shell with no elec- trons of higher energy causes it to lose an electron relatively easily. Sodium (Na) has a single electron in its 3satomic orbital. Consequently,sodium also has a relatively low ionization energy because, when it loses an electron, it is left with an outer shell of eight electrons. Elements (such as lithium and sodium) that have low ionization ener- gies are said to be electropositive—they readily lose an electron and thereby become positively charged. The elements in the first column of the periodic table are all electropositive—each readily loses an electron because each has a single electron in its outermost shell. Electrons in inner shells (those below the outermost shell) are called core electrons. Core electrons do not participate in chemical bonding. Electrons in the outermost shell are called valence electrons, and the outermost shell is called the valence shell. Car- bon,for example,has two core electrons and four valence electrons (Table 1.2). BRUI01-001_059r4 20-03-2003 2:58 PM Page 8 8 CHAPTER 1 Electronic Structure and Bonding • Acids and Bases Lithium and sodium each have one valence electron. Elements in the same column of the periodic table have the same number of valence electrons,and because the num- ber of valence electrons is the major factor determining an element’s chemical proper- ties, elements in the same column of the periodic table have similar chemical properties. Thus, the chemical behavior of an element depends on its electronic configuration. PROBLEM 4 Compare the ground-state electronic configurations of the following atoms,and check the relative positions of the atomsin Table1.3on p.10. a. carbon and silicon c. fluorine and bromine b. oxygen and sulfur d. magnesium and calcium When we draw the electrons around an atom, as in the following equations, core electrons are not shown; only valence electrons are shown. Each valence electron is shown as a dot. Notice that when the single valence electron of lithium or sodium is removed,the resulting atom—now called an ion—carries a positive charge. Li Li+ + e− Na Na+ + e− Fluorine has seven valence electrons (Table 1.2). Consequently,it readily acquires an electron in order to have an outer shell of eight electrons. When an atom acquires an electron, energy is released. Elements in the same column as fluorine (e.g., chlorine, bromine, and iodine) also need only one electron to have an outer shell of eight, so they,too,readily acquire an electron. Elements that readily acquire an electron are said to be electronegative—they acquire an electron easily and thereby become negatively charged. F + e− F − Cl + e− Cl − Ionic Bonds Because sodium gives up an electron easily and chlorine acquires an electron readily, when sodium metal and chlorine gas are mixed, each sodium atom transfers an elec- tron to a chlorine atom,and crystalline sodium chloride (table salt) is formed as a re- sult. The positively charged sodium ions and negatively charged chloride ions are 3-D Molecule: independent species held together by the attraction of opposite charges (Figure1.1). A Sodium chloride lattice bond is an attractive force between two atoms. Attractive forces between opposite charges are called electrostatic attractions. A bondthat is the result of only electro- static attractions is called an ionic bond. Thus,an ionic bondis formed when there is a transfer of electrons, causing one atom to become a positively charged ion and the other to become a negatively charged ion. Figure 1.1 N a. b. (a)Crystalline sodium chloride. (b)The electron-rich chloride ions are red and the electron-poor sodium ions are blue. Each chloride ion is surrounded by six sodium ions, and each sodium ion is surrounded by six chloride ions. Ingore the “bonds” holding the balls together; they are there only to keep the model from falling apart. BRUI01-001_059r4 20-03-2003 2:58 PM Page 9 Section 1.3 Ionic, Covalent, and Polar Bonds 9 ionic bond − + − Cl Na Cl + − + Na Cl Na − + − Cl Na Cl sodium chloride Sodium chloride is an example of an ionic compound. Ionic compounds are formed when an element on the left side of the periodic table (an electropositive ele- ment) transfers one or more electrons to an element on the right side of the periodic table (an electronegative element). Covalent Bonds Instead of giving up or acquiring electrons,an atom can achieve a filled outer shell by sharing electrons. For example, two fluorine atoms can each attain a filled shell of eight electrons by sharing their unpaired valence electrons. A bond formed as a result of sharing electronsis called a covalent bond. a covalent bond F + F F F Two hydrogen atoms can form a covalent bond by sharing electrons. As a result of co- valent bonding,each hydrogen acquires a stable,filled outer shell (with two electrons). H + H H H Similarly,hydrogen and chlorine can form a covalent bond by sharing electrons. In doing so,hydrogen fills its only shell and chlorine achieves an outer shell of eight electrons. H + Cl H Cl A hydrogen atom can achieve a completely empty shell by losing an electron. Loss of its sole electron results in a positively charged hydrogen ion. A positively charged hydrogen ion is called a protonbecause when a hydrogen atom loses its valence elec- tron, only the hydrogen nucleus—which consists of a single proton—remains. A hy- drogen atom can achieve a filled outer shell by gaining an electron,thereby forming a negatively charged hydrogen ion,called a hydride ion. H H+ + e– a hydrogen atom a proton H + e– H − a hydrogen atom a hydride ion Because oxygen has six valence electrons, it needs to form two covalent bonds to achieve an outer shell of eight electrons. Nitrogen, with five valence electrons, must Shown is a bronze sculpture of form three covalent bonds,and carbon,with four valence electrons,must form four co- Albert Einsteinon the grounds of valent bonds to achieve a filled outer shell. Notice that all the atoms in water,ammo- the National Academy of Sciences in nia,and methane have filled outer shells. Washington,DC. The statue mea- sures 21 feet from the top of the head 2 H + O H O to the tip of the feet and weighs 7000 H pounds. In his left hand,Einstein water holds the mathematical equations 3 H + N H N H that represent his three most impor- tant contributions to science:the H photoelectric effect,the equivalency ammonia of energy and matter,and the theory H of relativity. At his feet is a map of 4 H + C H C H the sky. H methane BRUI01-001_059r4 20-03-2003 2:58 PM Page 10 10 CHAPTER 1 Electronic Structure and Bonding • Acids and Bases Polar Covalent Bonds In the F¬F and H¬H covalent bonds shown previously, the atoms that share the bonding electrons are identical. Therefore, they share the electrons equally; that is, each electron spends as much time in the vicinity of one atom as in the other. An even (nonpolar) distribution of charge results. Such a bond is called a nonpolar covalent bond. In contrast, the bonding electrons in hydrogen chloride, water, and ammonia are more attracted to one atom than another because the atoms that share the electrons in these molecules are different and have different electronegativities. Electronegativity is the tendency of an atom to pull bonding electrons toward itself. The bonding elec- trons in hydrogen chloride, water, and ammonia molecules are more attracted to the atom with the greater electronegativity. This results in a polar distribution of charge. A polar covalent bondis a covalent bond between atoms of different electronegativities. The electronegativities of some of the elements are shown in Table 1.3. Notice that electronegativity increases as you go from left to right across a row of the periodic table or up any of the columns. A polar covalent bond has a slight positive charge on one end and a slight nega- tive charge on the other. Polarity in a covalent bond is indicated by the symbols d+ and d-,which denote partial positive and partial negative charges,respectively. The negative end of the bond is the end that has the more electronegative atom. The greater the difference in electronegativity between the bonded atoms,the more polar the bond will be. δ+ δ− δ+ δ− δ+ δ− δ+ H Cl H O H N H H H δ+ δ+ The direction of bond polarity can be indicated with an arrow. By convention, the arrow points in the direction in which the electrons are pulled,so the head of the arrow is at the negative end of the bond; a short perpendicular line near the tail of the arrow marks the positive end of the bond. H Cl TABLE 1.3 The Electronegativities of Selected Elementsa IA IIA IB IIB IIIA IVA VA VIA VIIA H 2.1 y t Li Be B C N O F vi 1.0 1.5 2.0 2.5 3.0 3.5 4.0 ati g e Na Mg Al Si P S Cl n o 0.9 1.2 1.5 1.8 2.1 2.5 3.0 tr c e K Ca Br el 0.8 1.0 2.8 g n I asi increasing electronegativity e 2.5 cr n i aElectronegativity values are relative, not absolute. As a result, there are several scales of electronegativities. The electronegativities listed here are from the scale devised by Linus Pauling.