1-2 Principles of Atomic Structure 5 directions of axe ucleu (comes out toward us) Figure 1-4 The are three 2p the 2porital each other.Each is labeled according on ong the orbital can hold a 2 ele 2p orbitals)can accommodate 8 electrons.and the third shell (one 3s orbital.three 3p orbitals.and five 3d orbitals)can accommodate 18 electrons. 1-2C Electronic Configurations of Atoms ctronic co e)stat GU th groun I we ave ed the p elements in the first two rows of the periodic table. TABLE 1-1 Electronic Configurations of the Elements of the First and Second Rows Element Configuration Valence Electrons Relative orbital energies 12.1 1292 一2p-2-2n 1s22s22p1 1s 2s energy -2s 2 1 56- Ne 1s2222p22p22p
1-2 Principles of Atomic Structure 5 « Figure 1-4 The 2p orbitals. There are three 2p orbitals, oriented at right angles to each other. Each is labeled according to its orientation along the x, y, or z axis. nucleus distance from the nucleus electron density node 2p the 2 px orbital the 2px, 2py, and 2pz orbitals superimposed at 90 angles x y z directions of axes (z comes out toward us) nodal plane x y x z y z The Pauli exclusion principle tells us that each orbital can hold a maximum of 2 electrons, provided that their spins are paired. The first shell (one 1s orbital) can accommodate 2 electrons. The second shell (one 2s orbital and three 2p orbitals) can accommodate 8 electrons, and the third shell (one 3s orbital, three 3p orbitals, and five 3d orbitals) can accommodate 18 electrons. 1-2C Electronic Configurations of Atoms Aufbau means “building up” in German, and the aufbau principle tells us how to build up the electronic configuration of an atom’s ground (most stable) state. Starting with the lowest-energy orbital, we fill the orbitals in order until we have added the proper number of electrons. Table 1-1 shows the ground-state electronic configurations of the elements in the first two rows of the periodic table. TABLE 1-1 Electronic Configurations of the Elements of the First and Second Rows Element Configuration Valence Electrons H 1 He 2 Li 1 Be 2 B 3 C 4 N 5 O 6 F 7 Ne 1s 8 2 2s2 2px 2 2py 2 2pz 2 1s2 2s2 2px 2 2py 2 2pz 1 1s2 2s2 2px 2 2py 1 2pz 1 1s2 2s2 2px 1 2py 1 2pz 1 1s2 2s2 2px 1 2py 1 1s2 2s2 2px 1 1s2 2s2 1s2 2s1 1s2 1s1 Relative orbital energies energy 2py 2pz 2px 1s 2s WADEMC01_0131478710.QXD 11/8/04 7:56 AM Page 5
6 Chapter 1:Introduction and Review Figure 1-5 Partial periodic table noble The organizat H A IIIA IVA VA VIA VIIA He Li Be B CNOF Ne AISi P S CI Ar Two additional concepts are illustrated in Table 1-1.The valence electrons are those electrons that are in the outermost shell.Carbon has four valence elec- rons,nitrogen has five,and oxygen has six.Helium has two valence electrons,and neon ha eight,corresponding to a fille shell and second shell, ctively.In general (for the repr ntative elements column or group num ne per electro nd th poth in the IA)of the periodic table.Carbon has four valence electrons,and it is in group IVA of the periodic table. Notice in Table 1-1 that carbon's third and fourth valence electrons are not red:they occupy separate orbitals.Although the Pasays two electron e orbita the airing of the energy md s whe e same orbital.The first 2p elect ne n orhital the last 2p orbital.The fourth.fifth.and sixth 2p electrons must pair up with the first three electrons. PROBLEM 1-1 ic configurations of the third-row elements shown in the partial periodic 1-3 Bond Formation: The Octet Rule the elect of a noble gas.such as He.Ne.or Ar.This principle has come to be called the octet rule because a filled shell implies eight valence electrons for the elements in the second row of the periodic table. 1-3A lonic Bonding as configurations.Some a .lithas co rations by transfe rom o less than the asily its valence electron.and fluorine easily gains one:
6 Chapter 1: Introduction and Review Lithium carbonate, a salt of lithium, is a mood-stabilizing agent used to treat the psychiatric disorder known as mania. Mania is characterized by behaviors such as elated mood, feelings of greatness, racing thoughts, and an inability to sleep. We don’t know how lithium carbonate helps to stabilize these patients’ moods. » Figure 1-5 First three rows of the periodic table. The organization of the periodic table results from the filling of atomic orbitals in order of increasing energy. For these representative elements, the number of the column corresponds to the number of valence electrons. IA H noble gases (VIII) Li Na IIA Be Mg IIIA B Al IVA C Si VA N P VIA O S VIIA F Cl He Ne Ar Partial periodic table 1-3 Bond Formation: The Octet Rule In 1915, G. N. Lewis proposed several new theories describing how atoms bond together to form molecules. One of these theories states that a filled shell of electrons is especially stable, and atoms transfer or share electrons in such a way as to attain a filled shell of electrons. A filled shell of electrons is simply the electron configuration of a noble gas, such as He, Ne, or Ar. This principle has come to be called the octet rule because a filled shell implies eight valence electrons for the elements in the second row of the periodic table. 1-3A Ionic Bonding There are two ways that atoms can interact to attain noble-gas configurations. Sometimes atoms attain noble-gas configurations by transferring electrons from one atom to another. For example, lithium has one electron more than the helium configuration, and fluorine has one electron less than the neon configuration. Lithium easily loses its valence electron, and fluorine easily gains one: Two additional concepts are illustrated in Table 1-1. The valence electrons are those electrons that are in the outermost shell. Carbon has four valence electrons, nitrogen has five, and oxygen has six. Helium has two valence electrons, and neon has eight, corresponding to a filled first shell and second shell, respectively. In general (for the representative elements), the column or group number of the periodic table corresponds to the number of valence electrons (Figure 1-5). Hydrogen and lithium have one valence electron, and they are both in the first column (group IA) of the periodic table. Carbon has four valence electrons, and it is in group IVA of the periodic table. Notice in Table 1-1 that carbon’s third and fourth valence electrons are not paired; they occupy separate orbitals. Although the Pauli exclusion principle says that two electrons can occupy the same orbital, the electrons repel each other, and pairing requires additional energy. Hund’s rule states that when there are two or more orbitals of the same energy, electrons will go into different orbitals rather than pair up in the same orbital. The first 2p electron (boron) goes into one 2p orbital, the second (carbon) goes into a different orbital, and the third (nitrogen) occupies the last 2p orbital. The fourth, fifth, and sixth 2p electrons must pair up with the first three electrons. PROBLEM 1-1 Write the electronic configurations of the third-row elements shown in the partial periodic table in Figure 1-5. WADEMC01_0131478710.QXD 11/8/04 7:56 AM Page 6
1-4 Lewis Structures 7 Li一f: Lit :F: >i计：f: electron transfer He configuration Ne configuration ionic bond A transfer of one electron gives each of these two elements a noble-gas config ration.The resulting ions have opposite charg and they attract each other to form an ionic bond.Ionic bonding usually results in the formation of a large crystal lattice rather than individual molecules.Ionic bonding is common in inorganic compounds but relatively uncommon in organic compounds 1-3B Covalent Bonding Covalent bonding come together and form a bond,they"share"their two electrons. and each atom has two electrons in its valence shell. H.+H. H:H each H shares two electrons (He configuration) We will study covalent bonding in more detail in Chapter 2. 1-4 n vale s sym ecroni yapaa by a d y o arrange Lewis Structures electrons s fo the sppro H :H or H-C-H H Carbon contributes four valence electrons,and each hydrogen contributes one,to give a total of eight electrons.All eight electrons surround carbon to give it an octet and HH HH ethane Once again,we have computed the total number of valence electrons(14)and distrib uted them so that each carbon atom is surrounded by 8 and each hydrogen by 2.The only possible structure for ethane is the one shown,with the two carbon atoms sharing ach hydrogen atom sharing a pair with one of the carbons.The eehemomoran haracterbstic of oofom strong carbon-car
1-4 Lewis Structures 7 Li F electron transfer Li He configuration F Ne configuration Li F ionic bond A transfer of one electron gives each of these two elements a noble-gas configuration. The resulting ions have opposite charges, and they attract each other to form an ionic bond. Ionic bonding usually results in the formation of a large crystal lattice rather than individual molecules. Ionic bonding is common in inorganic compounds but relatively uncommon in organic compounds. 1-3B Covalent Bonding Covalent bonding, in which electrons are shared rather than transferred, is the most common type of bonding in organic compounds. Hydrogen, for example, needs a second electron to achieve the noble-gas configuration of helium. If two hydrogen atoms come together and form a bond, they “share” their two electrons, and each atom has two electrons in its valence shell. We will study covalent bonding in more detail in Chapter 2. H H each H shares two electrons (He configuration) H H 1-4 Lewis Structures One way to symbolize the bonding in a covalent molecule is to use Lewis structures. In a Lewis structure, each valence electron is symbolized by a dot. A bonding pair of electrons is symbolized by a pair of dots or by a dash We try to arrange all the atoms so they have their appropriate noble-gas configurations: two electrons for hydrogen, and octets for the second-row elements. Consider the Lewis structure of methane Carbon contributes four valence electrons, and each hydrogen contributes one, to give a total of eight electrons. All eight electrons surround carbon to give it an octet, and each hydrogen atom shares two of the electrons. The Lewis structure for ethane is more complex. Once again, we have computed the total number of valence electrons (14) and distributed them so that each carbon atom is surrounded by 8 and each hydrogen by 2. The only possible structure for ethane is the one shown, with the two carbon atoms sharing a pair of electrons and each hydrogen atom sharing a pair with one of the carbons. The ethane structure shows the most important characteristic of carbon—its ability to form strong carbon–carbon bonds. H H C C H or 9C9C9H ethane H H H H H H H H 1C2H62 H H C H or 9C9H methane H H H H 1CH42. 1¬2. WADEMC01_0131478710.QXD 11/8/04 7:56 AM Page 7
8 Chapter 1:Introduction and Review Oxygen atoms,nitrogen atoms,and the halogens (F.Cl.Br.D usually have nonbonding electrons in their stable compounds.These lone pairs of nonbonding electrons help to determine the reactivity of their parent compounds.The following Lewis structures show one lone pair of electrons son the nitrogen atom of methy e pairs on t ogen atoms usually romethan H lone pair lone pairs H H N一H H-C-C-6: H-C 一Cl2 lone pairs HH HHH methylamine ethanol chloromethane A correct Lewis structure should show any lone pairs.Organic chemists often draw structures that omit most or all of the lone pairs.These are not true Lewis struc- tures because you must imagine the correct number of nonbonding electrons PROBLEM-SOLVING PROBLEM 1-2 Draw Lewis stn the way we rite (b)water.H2O help you throughout thiscourse. NH ()opt 2 (OH)CHs borane.BHs Explain what is unusual about the bonding in compounds in parts(1)and ( 1-5 drawing Lewis res in Sec ion 1-4.we placec t one pair Multiple Bonding nd.Many molecules have adiacent atoms sharing two or n three electron The sharing of two pairs is called a double bond,and the sharing of three pairs is called a triple bond.Ethylene (C2H4)is an organic compound with a double bond.When we draw a Lewis structure for ethylene,the only way to show both carbon atoms with octets is to draw them sh aring two pairs of electrons.The following ween two atoms s to give them octets.A doubled H. H or or or H CC-CH H H>C=0 H-C-N: H ethylene formaldehyde formaldimine e for parts of Acetvlene (CH)has a triple bond.Its Lewis stru re sho thre pipeline in Siberia. electrons ber een the carbon how oranic compounds with triple bonds.A triple dash( triple bond
8 Chapter 1: Introduction and Review Valence-shell electrons that are not shared between two atoms are called nonbonding electrons. A pair of nonbonding electrons is often called a lone pair. Oxygen atoms, nitrogen atoms, and the halogens (F, Cl, Br, I) usually have nonbonding electrons in their stable compounds. These lone pairs of nonbonding electrons help to determine the reactivity of their parent compounds. The following Lewis structures show one lone pair of electrons on the nitrogen atom of methylamine and two lone pairs on the oxygen atom of ethanol. Halogen atoms usually have three lone pairs, as shown in the structure of chloromethane. A correct Lewis structure should show any lone pairs. Organic chemists often draw structures that omit most or all of the lone pairs. These are not true Lewis structures because you must imagine the correct number of nonbonding electrons. PROBLEM 1-2 Draw Lewis structures for the following compounds. (a) ammonia, (b) water, (c) hydronium ion, (d) propane, (e) ethylamine, (f) dimethyl ether, (g) fluoroethane, (h) 2-propanol, (i) borane, (j) boron trifluoride, Explain what is unusual about the bonding in compounds in parts (i) and (j). BH3 BF3 CH3CH2F CH3CH1OH2CH3 CH3CH2NH2 CH3OCH3 H3O C3H8 + NH3 H2O H C N H H H H lone pair methylamine H C H H C H H O H lone pairs H C H H Cl ethanol chloromethane lone pairs 1-5 Multiple Bonding In drawing Lewis structures in Section 1-4, we placed just one pair of electrons between any two atoms. The sharing of one pair between two atoms is called a single bond. Many molecules have adjacent atoms sharing two or even three electron pairs. The sharing of two pairs is called a double bond, and the sharing of three pairs is called a triple bond. Ethylene is an organic compound with a double bond. When we draw a Lewis structure for ethylene, the only way to show both carbon atoms with octets is to draw them sharing two pairs of electrons. The following examples show organic compounds with double bonds. In each case, four electrons (two pairs) are shared between two atoms to give them octets. A double dash symbolizes a double bond. Acetylene has a triple bond. Its Lewis structure shows three pairs of electrons between the carbon atoms to give them octets. The following examples show organic compounds with triple bonds. A triple dash symbolizes a triple bond. 1‚2 1C2H22 C C HH HH or C C H H H H ethylene C H H C H H O O or formaldehyde C H H C H H N N or formaldimine H H 1“2 1C2H42 Acetylene, in combination with oxygen, burns with an intense flame that has diverse applications. It can be used for welding parts of a bridge underwater and for repairing an oil pipeline in Siberia. PROBLEM-SOLVING Hint Lewis structures are the way we write organic chemistry. Learning now to draw them quickly and correctly will help you throughout this course. WADEMC01_0131478710.QXD 11/8/04 7:56 AM Page 8
1-6 Electronegativity and Bond Polarity 9 H:C:C:H H:C:C:C:C:H H:C:C:N: or or or H H H一C=C-H H-C-C=C-C-H H-C-C=N: H acetylene dimethylacetylene All these Lewis structures show that carbon normally forms four bonds in neutral or ganic compounds.Nitrogen generally forms three bonds,and oxygen usually forms two.Hy drogen and the halogens usually form only one bond.The number of bonds an atom usually forms is cal alence.Ca nt,oxygen is d ing the of bonds fo ih If SUMMARY Common Bonding Patterns(Uncharged) -c- -N- -6- -H -c carbon nitrogen hydrogen halogens 3 PROBLEM 1-3 Write Lewis structures for the PROBLEM-SOLVING乙 (a)Na (d)co, (e)H2CNH (间C3H6 (two double bonds) nitroge single bond PRORLEM 1-4 bond,o Circle any lone pairs (pairs of nonbondingelectrons)in the structures you drew for Problem 1-3. hrtki9problems.consider A bond with the electrons shared equally between the two atoms is called a nonpolar 1-6 covalent bond.The bond in H2 and the C-C bond in ethane are nonpolar covalent bonds.In most bonds between two different elements.the bonding electrons are at- Electronegativity e.rrfand Bond Polarity H○H Na+:CI covalent bond ionic bond When carbon is bonded tochlorine.for example.the bonding electrons are attracted negativ 1-6 shows th ond polarity by an
1-6 Electronegativity and Bond Polarity 9 SUMMARY Common Bonding Patterns (Uncharged) carbon 4 0 valence: lone pairs: nitrogen 3 1 oxygen 2 2 halogens 1 3 hydrogen 1 0 9C9 9N9 9O9 9H 9Cl All these Lewis structures show that carbon normally forms four bonds in neutral organic compounds. Nitrogen generally forms three bonds, and oxygen usually forms two. Hydrogen and the halogens usually form only one bond. The number of bonds an atom usually forms is called its valence. Carbon is tetravalent, nitrogen is trivalent, oxygen is divalent, and hydrogen and the halogens are monovalent. By remembering the usual number of bonds for these common elements, we can write organic structures more easily. If we draw a structure with each atom having its usual number of bonds, a correct Lewis structure usually results. H9C#C9H H9C9C#C9C9H H H H9C9C#N H H H H acetylene dimethylacetylene acetonitrile H C C H H C C CC H H H C CH N H H or or or H H 1-6 Electronegativity and Bond Polarity A bond with the electrons shared equally between the two atoms is called a nonpolar covalent bond. The bond in and the bond in ethane are nonpolar covalent bonds. In most bonds between two different elements, the bonding electrons are attracted more strongly to one of the two nuclei. An unequally shared pair of bonding electrons is called a polar covalent bond. When carbon is bonded to chlorine, for example, the bonding electrons are attracted more strongly to the chlorine atom. The carbon atom bears a small partial positive charge, and the chlorine atom bears an equal amount of negative charge. Figure 1-6 shows the polar carbon–chlorine bond in chloromethane. We symbolize the bond polarity by an H ClC nonpolar covalent bond polar covalent bond ionic bond H C Na l + − H2 C¬C PROBLEM 1-3 Write Lewis structures for the following molecular formulas. (a) (b) HCN (c) HONO (d) (e) (f) (g) (h) HNNH (i) (j) (two double bonds) (k) (one triple bond) PROBLEM 1-4 Circle any lone pairs (pairs of nonbonding electrons) in the structures you drew for Problem 1-3. C3H4 C3H4 C2H3Cl C3H6 CO2 H2CNH HCO2H N2 PROBLEM-SOLVING Hint These “usual numbers of bonds” might be single bonds, or they might be combined into double and triple bonds. For example, three bonds to nitrogen might be three single bonds, one single bond and one double bond, or one triple bond In working problems, consider all possibilities. 1≠N‚N≠2. WADEMC01_0131478710.QXD 11/8/04 7:56 AM Page 9