1.2 ATOMIC STRUCTURE:ORBITALS 5 an orbital represents the space where an electron spends most Wha look like?There are four different kinds of orbitals denot rent shape.Of four,we'll be con pmd biolog rbitals orbital is oll-sha cloverleaf-sha wn in Figure 1.4.The fifth d orbital is shar ped like ar elongated dumbbell with a doughnut around its middle FIGURE 1.4 Representations of sp,and dorbitals.Ansorbital is sphe cl,ap orbita be Different lobes ofn orbitals are often drawn for convenience as hatofiore teardrops,but the ape An s orbital A p orbital A d orbital The orbitals in an atom are organized into different layers,or electron shells. of successively larger size and energy.Different shells contain different num bers and kinds of orbitals,and each orbital within a shell can be occupied by two electrons.The first shell contains only a single s orbital,denoted 1s,and thus holds only 2 electrons.The second shell contains one 2s orbital and three 2p orbitals and thus holds a total of 8 electrons.The third shell contains a 3s orbital,three 3p orbitals,and five 3d orbitals,for a total capacity of 18 electrons. These orbital groupings and their energy levels are shown in Figure 1.5. 3rd shell 3d (capacity-18electrons) shell holds a maximum of2 3s 十 electrons in one is orbital;the 2nd shell 2D (capacity-8 electrons) olds 15 one 3s,three 3p,and five3d orbit. n e are rep nergy level of the 4s orbital falls along mutua denoted px,Py and p.As shown in between 3p and 3d. "igure 1.6,the each p orbital are separa a region of zerd on two orbital regions by the no 18n th .cf in S ction 1 consequence t to
by saying that an orbital represents the space where an electron spends most (90%–95%) of its time. What do orbitals look like? There are four different kinds of orbitals, denoted s, p, d, and f, each with a different shape. Of the four, we’ll be concerned primarily with s and p orbitals because these are the most common in organic and biological chemistry. An s orbital is spherical, with the nucleus at its center; a p orbital is dumbbell-shaped; and four of the five d orbitals are cloverleaf-shaped, as shown in Figure 1.4. The fifth d orbital is shaped like an elongated dumbbell with a doughnut around its middle. An s orbital A p orbital A d orbital The orbitals in an atom are organized into different layers, or electron shells, of successively larger size and energy. Different shells contain different numbers and kinds of orbitals, and each orbital within a shell can be occupied by two electrons. The first shell contains only a single s orbital, denoted 1s, and thus holds only 2 electrons. The second shell contains one 2s orbital and three 2p orbitals and thus holds a total of 8 electrons. The third shell contains a 3s orbital, three 3p orbitals, and five 3d orbitals, for a total capacity of 18 electrons. These orbital groupings and their energy levels are shown in Figure 1.5. 3rd shell (capacity—18 electrons) 2nd shell (capacity—8 electrons) 1st shell (capacity—2 electrons) 3 d 3 p 3 s 2 p 2 s 1 s Energy The three different p orbitals within a given shell are oriented in space along mutually perpendicular directions, denoted px, py, and pz. As shown in Figure 1.6, the two lobes of each p orbital are separated by a region of zero electron density called a node. Furthermore, the two orbital regions separated by the node have different algebraic signs, and , in the wave function, as represented by the different colors in Figure 1.6. As we’ll see in Section 1.11, the algebraic signs of the different orbital lobes have important consequences with respect to chemical bonding and chemical reactivity. FIGURE 1.4 Representations of s, p, and d orbitals. An s orbital is spherical, a p orbital is dumbbellshaped, and four of the five d orbitals are cloverleaf-shaped. Different lobes of p orbitals are often drawn for convenience as teardrops, but their true shape is more like that of a doorknob, as indicated. FIGURE 1.4 Representations of s, p, and d orbitals. An s orbital is spherical, a p orbital is dumbbellshaped, and four of the five d orbitals are cloverleaf-shaped. Different lobes of p orbitals are often drawn for convenience as teardrops, but their true shape is more like that of a doorknob, as indicated. FIGURE 1.5 The energy levels of electrons in an atom. The first shell holds a maximum of 2 electrons in one 1s orbital; the second shell holds a maximum of 8 electrons in one 2s and three 2p orbitals; the third shell holds a maximum of 18 electrons in one 3s, three 3p, and five 3d orbitals; and so on. The two electrons in each orbital are represented by up and down arrows, hg. Although not shown, the energy level of the 4s orbital falls between 3p and 3d. FIGURE 1.5 The energy levels of electrons in an atom. The first shell holds a maximum of 2 electrons in one 1s orbital; the second shell holds a maximum of 8 electrons in one 2s and three 2p orbitals; the third shell holds a maximum of 18 electrons in one 3s, three 3p, and five 3d orbitals; and so on. The two electrons in each orbital are represented by up and down arrows, hg. Although not shown, the energy level of the 4s orbital falls between 3p and 3d. 1.2 atomic structure: orbitals 5 39144_01_0001-0032.indd 5 7/27/09 1:28:35 PM
CHAPTER 1 STRUCTURE AND BONDING ree mu shaped orbitals has two lobes separated by a node.The two lobes have different algebraic different colors A2px orbital A2py orbital A2pz orbital 18 Atomic Structure:Electron Configurations The lowest-energy arrangement or ground-state electron configuration.of an atom is a listing of the orbitals oco upied by its electrons.We can predict this arrangement by following three rules: Rule 1 The lowest-energy orbitals fill up first,according to the order 1s-2s-2p 3s→3p→4s→3d,a statement called the aufbau principle.Note that the4s orbital lies between the 3p and 3d orbitals in energy. Rule Electrons act in some ways as if they were spinning around an axis,in much the same way that the earth spins.This spin can have two orientations,denoted as up f and down Only two electrons can occupy an orbital,and they must be of opposite spin,a statement called the Pauli exclusion principle mpty orbitals e,one vith spins parallel until examples of how th e rules are sh own in Table 1.1.Hydro cupy s,hydro d-state electrons a the 12262 1 nd so forth Noto that a ipt is used to repres nt the ons in a par- TABLE 1.1 Ground-State Electron Configurations of Some Elements Atomic Atomic Element number Configuration Element number Configuration Hydrogen 1 18十- Phosphorus 15 3p +++ 3s Carbon 2p ++ 2p 2s 1s 1s
6 chapter 1 structure and bonding A 2px orbital A 2py orbital y A 2pz orbital x z y x z y x z 1.3 Atomic Structure: Electron Configurations The lowest-energy arrangement, or ground-state electron configuration, of an atom is a listing of the orbitals occupied by its electrons. We can predict this arrangement by following three rules: Rule 1 The lowest-energy orbitals fill up first, according to the order 1s n 2s n 2p n 3s n 3p n 4s n 3d, a statement called the aufbau principle. Note that the 4s orbital lies between the 3p and 3d orbitals in energy. Rule 2 Electrons act in some ways as if they were spinning around an axis, in much the same way that the earth spins. This spin can have two orientations, denoted as up h and down g. Only two electrons can occupy an orbital, and they must be of opposite spin, a statement called the Pauli exclusion principle. Rule 3 If two or more empty orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half-full, a statement called Hund’s rule. Some examples of how these rules apply are shown in Table 1.1. Hydrogen, for instance, has only one electron, which must occupy the lowest-energy orbital. Thus, hydrogen has a 1s ground-state configuration. Carbon has six electrons and the ground-state configuration 1s2 2s2 2px 1 2py 1, and so forth. Note that a superscript is used to represent the number of electrons in a particular orbital. FIGURE 1.6 Shapes of the 2p orbitals. Each of the three mutually perpendicular, dumbbellshaped orbitals has two lobes separated by a node. The two lobes have different algebraic signs in the corresponding wave function, as indicated by the different colors. FIGURE 1.6 Shapes of the 2p orbitals. Each of the three mutually perpendicular, dumbbellshaped orbitals has two lobes separated by a node. The two lobes have different algebraic signs in the corresponding wave function, as indicated by the different colors. TABLE 1.1 Ground-State Electron Configurations of Some Elements Atomic Atomic Element number Confi guration Element number Confi guration Hydrogen 1 1 s Phosphorus 15 3 s 2 s 1 s 3 p Carbon 6 2 p 2 s 1 s 2 p 39144_01_0001-0032.indd 6 7/27/09 1:28:36 PM
1.4 DEVELOPMENT OF CHEMICAL BONDING THEORY 7 Problem1.1 Give the ground-state electron configuration for each of the following elements: (a)Oxygen (b)Phosphorus (c)Sulfur nanyolectroasdbesesl ach of the following biological trace elements have (a)Magnesium Cob t el balt (c)Selenium 14Development of Chemical Bonding Theory By the mid-1800s,the new science of chemistry was developing rapidly and chemists had hegun to probe the forces holding comnounds together in 1858 August Kekul6 and Archibald Couper independently proposed that,in all its compounds,carbon is tetravalent-it always forms four bonds when it joins other elements to form stable compounds.Furthermore,said Kekul6,carbon atoms can bond to one another to form extended chains of linked atoms. Shortly after the tetravalent nature of carbon was proposed,extensions to the Kekul6-Couper theory were made when the possibility of multiple bond- ing between atoms was suggested.Emil Erlenmeyer proposed a carbon-carbon triple bond or acetylene,and Alexander Crum Brown proposed a carbon- carbon double bond for ethylene.In 1865.Kekule provide another majo advance when he suggested that carbon chains can double back on themselves ug Kekuand Couper cribing the etravale Le Belac on hird unti ed co ey propos d e fou i s.Van't Hoff nt d ed that the four which is bonded sit at the c tetrahe carbon in the center A re resentation of a tetrahedral carbon atom is showr n in Figure 17.Note the conventions used to show three-dimensionality:solid lines bonds in the plane of the page,the heavy wedged line represents a bond com- ing out of the r will be used throughout this text. Bonds in plane eding of page atom the solid lines renresen bonds in the plane of the paper. he heavy wedged lir rep ent ona com ashed line behind the plane of the page
Problem 1.1 Give the ground-state electron configuration for each of the following elements: (a) Oxygen (b) Phosphorus (c) Sulfur Problem 1.2 How many electrons does each of the following biological trace elements have in its outermost electron shell? (a) Magnesium (b) Cobalt (c) Selenium 1.4 Development of Chemical Bonding Theory By the mid-1800s, the new science of chemistry was developing rapidly and chemists had begun to probe the forces holding compounds together. In 1858, August Kekulé and Archibald Couper independently proposed that, in all its compounds, carbon is tetravalent—it always forms four bonds when it joins other elements to form stable compounds. Furthermore, said Kekulé, carbon atoms can bond to one another to form extended chains of linked atoms. Shortly after the tetravalent nature of carbon was proposed, extensions to the Kekulé–Couper theory were made when the possibility of multiple bonding between atoms was suggested. Emil Erlenmeyer proposed a carbon–carbon triple bond for acetylene, and Alexander Crum Brown proposed a carbon– carbon double bond for ethylene. In 1865, Kekulé provided another major advance when he suggested that carbon chains can double back on themselves to form rings of atoms. Although Kekulé and Couper were correct in describing the tetravalent nature of carbon, chemistry was still viewed in a two-dimensional way until 1874. In that year, Jacobus van’t Hoff and Joseph Le Bel added a third dimension to our ideas about organic compounds. They proposed that the four bonds of carbon are not oriented randomly but have specific spatial directions. Van’t Hoff went even further and suggested that the four atoms to which carbon is bonded sit at the corners of a regular tetrahedron, with carbon in the center. A representation of a tetrahedral carbon atom is shown in Figure 1.7. Note the conventions used to show three-dimensionality: solid lines represent bonds in the plane of the page, the heavy wedged line represents a bond coming out of the page toward the viewer, and the dashed line represents a bond receding back behind the page away from the viewer. These representations will be used throughout this text. Bond receding into page Bonds in plane of page Bond coming out of plane A tetrahedral carbon atom A regular tetrahedron H H H H C FIGURE 1.7 A representation of van’t Hoff’s tetrahedral carbon atom. The solid lines represent bonds in the plane of the paper, the heavy wedged line represents a bond coming out of the plane of the page, and the dashed line represents a bond going back behind the plane of the page. FIGURE 1.7 A representation of van’t Hoff’s tetrahedral carbon atom. The solid lines represent bonds in the plane of the paper, the heavy wedged line represents a bond coming out of the plane of the page, and the dashed line represents a bond going back behind the plane of the page. 1.4 development of chemical bonding theory 7 39144_01_0001-0032.indd 7 7/27/09 1:28:37 PM
8 CHAPTER 1 STRUCTURE AND BONDING Why.though,do atoms bond together,and how can bonds be described electronically?The why question is relatively easy to answer:atoms bond together because the compound that results is more stable and lower in energy than the separate atoms.Energy(usually as heat)is always released and flows out of the chemical system when a chemical bond forms.Conversely.energy must be put into the system to break a chemical bond.Making bonds always releases energy,and breaking bonds always absorbs energy.The how question is more difficult.To answer it,we need to know more about the electronic properties of atoms. We know through observation that eight electrons(an electron octet)in an atom's outermost shell,or valence sh ility to the noble the period e(2 +8跳：Ar(2 8+8 also knov hat t emistry of n-group is gove est noble gas nguration of the near n group 12 r ex ac -ga ing the single sele valence h o up not rati nion. Th tan h nds like c1- by an electrostatic attrac tion that ts clo er to the middle of the riodic table form bonds?Lo k at methane.ch the main constituent of p ple.Tho bonding in mothane is not ionic because it would take toomuch noSoteabonas22s2p918n gain or lose four electrons to achieve ration As a result carbon honds to other atoms not hy electrons,but by sharing them.Such a shared-electron posed in 1916 by G.N.Lewis,is called a covalent bond. The neutral collection of atoms held together by covalent bonds is called a molecule. A simple way of indicating the covalent bonds in molecules is to use what are called Lewis structures,or electron-dot structures,in which the valence shell electrons of an atom are represented as dots.Thus,hydrogen has one dot representing its 1selectron,carbon has four dots(2s2 2p2),oxygen has six dots (2s2 2p4).and so on.A stable molecule results whenever a noble-gas configu- ration is achieved for all the atoms- eight dots (an octet)for main-group at ms drawn between atoms H:O:H H-C-H H-N-H H-6-H H-C-6-H Atnia 0O
8 chapter 1 structure and bonding Why, though, do atoms bond together, and how can bonds be described electronically? The why question is relatively easy to answer: atoms bond together because the compound that results is more stable and lower in energy than the separate atoms. Energy (usually as heat) is always released and flows out of the chemical system when a chemical bond forms. Conversely, energy must be put into the system to break a chemical bond. Making bonds always releases energy, and breaking bonds always absorbs energy. The how question is more difficult. To answer it, we need to know more about the electronic properties of atoms. We know through observation that eight electrons (an electron octet) in an atom’s outermost shell, or valence shell, impart special stability to the noblegas elements in group 8A of the periodic table: Ne (2 8); Ar (2 8 8); Kr (2 8 18 8). We also know that the chemistry of main-group elements is governed by their tendency to take on the electron configuration of the nearest noble gas. The alkali metals in group 1A, for example, achieve a noble-gas configuration by losing the single s electron from their valence shell to form a cation, while the halogens in group 7A achieve a noble-gas configuration by gaining a p electron to fill their valence shell and form an anion. The resultant ions are held together in compounds like Na Cl by an electrostatic attraction that we call an ionic bond. But how do elements closer to the middle of the periodic table form bonds? Look at methane, CH4, the main constituent of natural gas, for example. The bonding in methane is not ionic because it would take too much energy for carbon (1s2 2s2 2p2) to either gain or lose four electrons to achieve a noble-gas configuration. As a result, carbon bonds to other atoms, not by gaining or losing electrons, but by sharing them. Such a shared-electron bond, first proposed in 1916 by G. N. Lewis, is called a covalent bond. The neutral collection of atoms held together by covalent bonds is called a molecule. A simple way of indicating the covalent bonds in molecules is to use what are called Lewis structures, or electron-dot structures, in which the valenceshell electrons of an atom are represented as dots. Thus, hydrogen has one dot representing its 1s electron, carbon has four dots (2s2 2p2), oxygen has six dots (2s2 2p4), and so on. A stable molecule results whenever a noble-gas configuration is achieved for all the atoms—eight dots (an octet) for main-group atoms or two dots for hydrogen. Simpler still is the use of Kekulé structures, or linebond structures, in which a two-electron covalent bond is indicated as a line drawn between atoms. C HH H H H C H H N HH H O HH O H C HH H H N HH H H O Water (H2O) H H C H H Methane (CH4) Electron-dot structures (Lewis structures) Line-bond structures (Kekulé structures) Ammonia (NH3) Methanol (CH3OH) O H 39144_01_0001-0032.indd 8 7/27/09 1:28:37 PM
1.4 DEVELOPMENT OF CHEMICAL BONDING THEORY 9 s on how many addi Ce e ctrons(1s o rea a noble-gas co ch the heliu Hydroger ron so it forms one 10u 22 de Ni and forms two bond and t have 2s22D4). en vale ace electrons.need one m e.and form one bond. c :F-:CI- H一 -N- -6- 8r-i- One bond Four bonds Three bonds Two bonds One bond Valence electrons that are not used for bonding are called lone-pair electrons,or nonbonding electrons.The nitrogen atom in ammonia(NH3) for instance,shares six valence electrons in three covalent bonds and has its ence electrons in a nonbonding lone pair.As a time-saving nonbondi awing line-bond nave to keep them in mind since they're often cru- reactions Nonbonding lone-pair electrons Ammonia WORKED EXAMPLE.Predicing the Number of Bonds Formed by Atoms in a Molecule How many hydrogen atoms does phosphorus bond to in phosphine,PH? Strategy Identify the periodic group of phosphorus,and tell from that how many elec- trons (bonds)are needed to make an octet. Solution Phosphorus,like nitrogen,is in group 5A of the periodic table and has five vale ce electrons.It thus needs to share three more electrons to make an octet and therefore bonds to three hydrogen atoms,giving PH3. lecule of chloroform,CHCl3,using solid,wedged,and dashed lines show its tetra edral geometry
The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble-gas configuration. Hydrogen has one valence electron (1s) and needs one more to reach the helium configuration (1s2), so it forms one bond. Carbon has four valence electrons (2s2 2p2) and needs four more to reach the neon configuration (2s2 2p6), so it forms four bonds. Nitrogen has five valence electrons (2s2 2p3), needs three more, and forms three bonds; oxygen has six valence electrons (2s2 2p4), needs two more, and forms two bonds; and the halogens have seven valence electrons, need one more, and form one bond. One bond Four bonds Three bonds Two bonds One bond Br F Cl I H C N O Valence electrons that are not used for bonding are called lone-pair electrons, or nonbonding electrons. The nitrogen atom in ammonia (NH3), for instance, shares six valence electrons in three covalent bonds and has its remaining two valence electrons in a nonbonding lone pair. As a time-saving shorthand, nonbonding electrons are often omitted when drawing line-bond structures, but you still have to keep them in mind since they’re often crucial in chemical reactions. Nonbonding, lone-pair electrons H N H H or H N H or H H N H H Ammonia WORKED EXAMPLE 1.1 Predicting the Number of Bonds Formed by Atoms in a Molecule How many hydrogen atoms does phosphorus bond to in phosphine, PH?? Strategy Identify the periodic group of phosphorus, and tell from that how many electrons (bonds) are needed to make an octet. Solution Phosphorus, like nitrogen, is in group 5A of the periodic table and has five valence electrons. It thus needs to share three more electrons to make an octet and therefore bonds to three hydrogen atoms, giving PH3. Problem 1.3 Draw a molecule of chloroform, CHCl3, using solid, wedged, and dashed lines to show its tetrahedral geometry. 1.4 development of chemical bonding theory 9 39144_01_0001-0032.indd 9 7/27/09 1:28:37 PM