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Capítulo 1 Átomos: O Mundo Quântico 147 Os elementos do Período 5 adicionam outros 18 elétrons, com o preenchimento dos or- bitais 5s, 4d e 5p. No Período 6, mais 32 elétrons são adicionados, porque também são adi- cionados os 14 elétrons dos sete orbitais 4f. Os elementos do bloco f têm propriedades quí- micas muito semelhantes, porque sua configuração eletrônica difere somente na população dos orbitais f internos e estes elétrons participam pouco da formação de ligação. Os blocos da Tabela Periódica são nomeados segundo o último orbital que é ocupado de acordo com o princípio da construção. Os períodos são numerados de acordo com o número quântico principal da camada de valência. A PERIODICIDADE DAS PROPRIEDADES DOS ÁTOMOS A Tabela Periódica pode ser usada na previsão de muitas propriedades, muitas das quais são cruciais para a compreensão dos materiais (Seções 1.20 e 1.21) e das ligações químicas (Capí- tulos 2 e 3), e para a organização dos elementos de acordo com essas propriedades (Capítulos 14 a 16). A variação da carga nuclear efetiva na Tabela Periódica tem papel importante na ex- plicação das tendências da periodicidade. A Figura 1.40 mostra a variação da carga efetiva nos três primeiros períodos. Ela cresce da esquerda para a direita em cada período e cai rapidamen- te na passagem de um período para o outro. 1.14 Raio Atômico As nuvens de elétrons não têm fronteiras bem definidas; logo, não é possível medir o raio exato de um átomo. Entretanto, quando os átomos se organizam como sólidos e moléculas, seus cen- tros encontram-se em distâncias definidas uns dos outros. O raio atômico de um elemento é de- finido como sendo a metade da distância entre os núcleos de átomos vizinhos (11). Se o elemen- to é um metal, o raio atômico é a metade da distância entre os centros de átomos vizinhos em uma amostra sólida. Por exemplo, como a distância entre os núcleos vizinhos do cobre sólido é 256 pm, o raio atômico do cobre é 128 pm. Se o elemento é um não-metal ou um metalóide, usamos a distância entre os núcleos de átomos unidos por uma ligação química. Esse raio é tam- bém chamado de raio covalente do elemento. Como exemplo, a distância entre os núcleos de uma molécula de Cl2 é 198 pm; logo, o raio covalente do cloro é 99 pm. Se o elemento é um gás nobre, nós usamos o raio de van der Waals, que é a metade da distância entre os centros de áto- mos vizinhos em uma amostra do gás sólido. Os raios atômicos dos gases nobres listados no Apêndice 2 são todos raios de van der Waals. Como os átomos de uma amostra de gás nobre não estão ligados quimicamente, os raios de van der Waals são, em geral, muito maiores do que os raios covalentes, e é melhor não incluí-los em nossa discussão das tendências de periodicidade. A Figura 1.41 mostra alguns raios atômicos e a Figura 1.42 mostra a variação do raio atô- mico com o número atômico. Observe a periodicidade, isto é, o padrão dentado na segunda. O raio atômico geralmente decresce da esquerda para a direita ao longo de um período e cresce com o valor de n em cada grupo. C ar ga n uc le ar e fe tiv a, Z e f 1 2 4 5 6 7 Número atômico, Z 1 3 5 7 8 11 139 15 3 He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar FIGURA 1.40 Variação da carga nuclear efeti- va do elétron de valência mais externo com o número atômi- co. Observe que a carga nu- clear efetiva aumenta da es- querda para a direita no perío- do, mas cai quando o elétron mais externo ocupa uma nova camada. (A carga nuclear efeti- va é, na verdade, Zefe, porém Zef é comumente chamado de carga) As propriedades atômicas dos elementos estão listadas no Apêndice 2D Johannes van de Waals foi um cientista holandês que estudou as interações entre moléculas. Veja o Capítulo 4. 11 Raio iônico 2r Carga nuclear efe,va (Z*) x Número atômico (Z) ATKINS, P.; JONES, L. Princípios de química: ques2onando a vida moderna e o meio ambiente. 3.ed. Bookman, 2006. Energias em átomos mul,eletrônicos 4d Same n ! !, different n Same n, different ! n n ! !! 4 4 4 3 3 3 2 2 1 2 1 0 2 1 0 1 0 0 6 5 4 5 4 3 3 2 1 EN ER GY 4p 4s 3d 1s 3p 3s 2p 2s 340 Chapter 8 Atomic Electron Configurations and Chemical Periodicity Active Figure 8.4 Experimentally determined order of subshell energies. Energies of electron shells increase with increasing n and, within a shell, subshell energies increase with increasing . (The energy axis is not to scale.) The energy gaps between subshells of a given shell become smaller as n increases. Note that the order of orbital energies does not correspond to the order of orbital filling for the heavier elements. For the filling order, see Figure 8.5. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. / The experimentally determined order of subshell energies in Figure 8.4 shows that the subshell energies of multielectron atoms depend on both n and . The subshells with n" 3, for example, have different energies; for a given atom they are in the order 3s # 3p # 3d. The subshell energy order in Figure 8.4 and the actual electron arrangements of the elements lead to two general rules that help us predict these arrangements: • Electrons are assigned to subshells in order of increasing “n! ” value. • For two subshells with the same value of “n! ,” electrons are assigned first to the subshell of lower n. The following are examples of these rules. • Electrons are assigned to the 2s subshell (n! " 2! 0" 2) before the 2p sub- shell (n! " 2! 1" 3). • Electrons are assigned in the order 3s (n! " 3! 0" 3) before 3p (n! " 3! 1" 4) before 3d (n! " 3! 2" 5). • Electrons fill the 4s subshell (n! " 4) before filling the 3d subshell (n! " 5). These filling orders, summarized in Figure 8.5, have been amply verified by experi- ment. // / // / / / / / KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. As subcamadas são preenchidas na ordem das somas n+l crescentes Exercise 8.1—Order of Subshell Assignments Using the “n! ” rules, you can generally predict the order of subshell assignments (the electron filling order) for a multielectron atom. To which of the following subshells should an electron be assigned first? (a) 4s or 4p (b) 5d or 6s (c) 4f or 5s Effective Nuclear Charge, Z* The order in which electrons are assigned to subshells in an atom, and many atomic properties, can be rationalized by the concept of effective nuclear charge (Z*). This is the nuclear charge experienced by a particular electron in a multielectron atom, as modified by the presence of the other electrons. In the hydrogen atom, with only one electron, the 2s and 2p subshells have the same energy. However, for lithium, an atom with three electrons, the presence of the 1s electrons alters the energy of the 2s and 2p subshells. Why should this be true? This question can be answered in part by referring to Figure 8.6. Figure 8.6 plots, qualitatively, the surface density function (4 r 2 2) for a 2s elec- tron [! Figure 7.14]. The probability of finding the electron (vertical axis) changes as one moves away from the nucleus (horizontal axis). Lightly shaded on this figure is the region occupied by the 1s electrons of lithium. Observe that the 2s electron wave occurs partly within the region occupied by 1s electrons. Chemists say that the 2s elec- tron density region penetrates the 1s electron density region. This alters the energy of the 2s electronrelative to what it would be in the H atom where there are no other electrons. As more electrons are added to an atom, the outermost electrons will pen- etrate the region occupied by the inner electrons, but the penetration is different for ns, np, and nd orbitals, and their energies are altered by differing amounts. cp / 8.3 Atomic Subshell Energies and Electron Assignments 341 n value ! value 0 8 8s 7s 6s 5s 4s 3s 2s 7p 6p 5p 4p 3p 6d 5d n + ! = 2 n + ! = 1 n + ! = 3 n + ! = 4 n + ! = 5 n + ! = 6 n + ! = 7 n + ! = 8 4d 5f 4f 3d 2p 1s 7 6 5 4 3 2 1 1 2 3 Figure 8.5 Subshell filling order. Subshells in atoms are filled in order of increasing n! . When two subshells have the same n! value, the subshell of lower n is filled first. To use the diagram, begin at 1s and follow the arrows of increasing n! . (Thus, the order of filling is 1s 2s 2p 3s 3p 4s 3d and so on.) 11111 1/ / / ■ More About Z* For a more complete discussion of effective nuclear charge, see D. M. P. Mingos: Essential Trends in Inorganic Chemistry, New York, Oxford University Press, 1998. KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. Configurações eletrônicas e tabela periódica s–block elements p–block elements d–block elements (transition metals) f–block elements: lanthanides (4f) and actinides (5f) 1s 1s 3d 4d 5d 6d 2s 3s 4s 5s 6s 2p 3p 4p 5p 6p 4f 5f 7s urations, but the noble gas notation also conveys the idea that the core electrons can generally be ignored when considering the chemistry of an element. The electrons beyond the core electrons—the 2s1 electron in the case of lithium—are the valence electrons, the electrons that determine the chemical properties of an element. The position of lithium in the periodic table tells you its configuration immedi- ately. All the elements of Group 1A have one electron assigned to an s orbital of the nth shell, for which n is the number of the period in which the element is found (Figure 8.7). For example, potassium is the first element in the n! 4 row (the fourth period), so potassium has the electron configuration of the element preced- ing it in the table (Ar) plus a final electron assigned to the 4s orbital: [Ar]4s1. Beryllium (Be) and Other Elements of Group 2A Beryllium, in Group 2A, has two electrons in the 1s orbital plus two additional electrons. All elements of Group 2A have electron configurations of [electrons of preceding noble gas]ns2, where n is the period in which the element is found in the periodic table. Because all the elements of Group 1A have the valence electron configuration ns1, and those in Group 2A have ns2, these elements are called s -block elements. Boron (B) and Other Elements of Group 3A Boron (Group 3A) is the first element in the block of elements on the right side of the periodic table. Because the 1s and 2s orbitals are filled in a boron atom, the fifth electron must be assigned to a 2p orbital. Boron: spdf notation 1s22s22p1 Box notation 1s 2s 2p or [He]2s22p1 Beryllium: spdf notation 1s22s2 Box notation 1s 2s 2p or [He]2s2 8.4 Atomic Electron Configurations 345 Active Figure 8.7 Electron configurations and the periodic table. The outermost electrons of an element are assigned to the indicated orbitals. See Table 8.3. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. Raio atômico em pm para elementos do grupo principal 8.6 Atomic Properties and Periodic Trends 355 C C Cl Cl 154 pm 198 pm 176 pm (a) (b) C Cl A distance equivalent to 4 times the radius of an aluminum atom Figure 8.10 Determining atomic radii. (a) The sum of the atomic radii of C and Cl provides a good esti- mate of the C Cl distance in a molecule having such a bond. (b) Each sphere in this tiny piece of an alu- minum crystal represents an aluminum atom. Measuring the distance shown, for example, allows a scientist to estimate the radius of an aluminum atom. ¬ MAIN GROUP METALS TRANSITION METALS NONMETALS METALLOIDS 1A 1A 2A 3A 4A 5A 6A 7A H, 37 Li, 152 Na, 186 K, 227 Rb, 248 Cs, 265 Be, 113 Mg, 160 Ca, 197 Sr, 215 Ba, 217 B, 83 Al, 143 Ga, 122 In, 163 Tl, 170 C, 77 Si, 117 Ge, 123 Sn, 141 Pb, 154 N, 71 P, 115 As, 125 Sb, 141 Bi, 155 O, 66 S, 104 Se, 117 Te, 143 Po, 167 F, 71 Cl, 99 Br, 114 I, 133 Active Figure 8.11 Atomic radii in picometers for main group elements. 1 pm! 1" 10#12 m. Data taken from J. Emsley: The Elements, 3rd ed., Oxford, Clarendon Press, 1998. See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. Raios atômicos 356 Chapter 8 Atomic Electron Configurations and Chemical Periodicity 100 150 200 250 150 200 250 R ad iu s (p m ) Period 6th Period 5th Period 4th Period Cs Rb K Hg Cd Zn Au Ag Cu PtIrOsRe W Pd Zr Nb Mo Tc Ru Rh NiCoFeMn CrV Ti Sc Ca 1A 4 5 6 8B 2A 3B 4B 5B 6B 7B 1B 2B Transition metals Figure 8.12 Trends in atomic radii for the transition elements. Atomic radii of the Group 1A and 2A metals and the transition metals of the fourth, fifth, and sixth periods. • The size of an atom is determined by the outermost electrons. In going from the top to the bottom of a group in the periodic table, the outermost electrons are assigned to orbitals with increasingly higher values of the principal quan- tum number, n. The underlying electrons require some space, so the electrons of the outer shell must be farther from the nucleus. • For main group elements of a given period, the principal quantum number, n, of the valence electron orbitals is the same. In going from one element to the next across a period, a proton is added to each nucleus and an electron is added to each outer shell. In each step, the effective nuclear charge, Z* [! Table 8.2] increases slightly because the effect of each additional proton is more impor- tant than the effect of an additional electron. The result is that attraction be- tween the nucleus and electrons increases, and atomic radius decreases. The periodic trend in the atomic radii of transition metal atoms (Figure 8.12) is somewhat different from that for main group elements. Going from left to right across a given period, the radii initially decrease across the first few elements. The sizes of the elements in the middle of a transition series then change very little un- til a small increase in size occurs at the end of the series. The size of the atom is de- termined largely by electrons in the outermost shell—that is, by the electrons of the ns subshell. In the first transition series, for example, the outer shell contains the 4s electrons, but electrons are being added to the 3d orbitals across the series. The in- creased nuclear charge on the atoms as one moves from left to right should cause the radius to decrease. This effect, however, is mostly cancelled out by increased electron–electron repulsion among the electrons. On reaching the Groups 1B and 2B elements at the end of the series, the size increases slightly because the d subshell is filled, and electron–electron repulsions cause the size to increase.■ Trends in Atomic Radii General trends in atomic radii of s- and p-block elements with position in the periodic table. Increase Atomic radiiIn cr ea se KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. Energias de ionização As another example, consider the first two ionization energies for lithium. First ionization energy, IE1! 513.3 kJ/mol Li(g) Li"(g)" e# 1s22s1 1s2 Second ionization energy, IE2! 7298 kJ/mol Li"(g) Li2"(g) " e# 1s2 1s1 The ionization energy for the removal of the second electron is large because the second electron is removed from a much lower energy (inner) subshell. For main group (s- and p-block) elements, first ionization energies generally increase across a period and decrease down a group (Figure 8.13 and Appendix F). The trend across a period is rationalized by the increase in effective nuclear charge, Z*, with in- creasing atomic number. Not only does this mean that the atomic radius decreases, but the energy required to remove an electron also increases. The general decrease in ionization energy down a group occurs because the electron removed is increas- ingly farther from the nucleus, thus reducing the nucleus-electron attractive force. A closer look at ionization energies reveals that the trend across a given period is not smooth, particularly in the second period. Variations are seen on going from s-block to p-block elements—from beryllium to boron, for example. This occurs be- cause the 2p electrons are slightly higher in energy than the 2s electrons (see Figure 8.4, page 340), so the ionization energy for boron is lower than that for beryllium. Moving from boron to carbon and then to nitrogen, the effective nuclear charge increases (see Table 8.2), which again means an increase in ionization en- ergy. Another dip to lower ionization energy occurs on passing from Group 5A to Group 6A. This is especially noticeable in the second period (N and O). No change occurs in either n or , but electron–electron repulsions increase for the following/ ¡ ¡ 358 Chapter 8 Atomic Electron Configurations and Chemical Periodicity 0 500 1000 1500 2000 2500 1 2 3 4 Ne F O N C B Be Li H He Ar Cl SP Si Al Mg Fi rs t i on iz at io n en er gy (k J/ m ol ) 1A 2A 3A 4A 5A 6A 7A 8A Group Period Na Kr Br SeAs Ge GaCa K Active Figure 8.13 First ionization energies of the main group elements of the first four periods. (For data on all the elements see Appendix F.) See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accom- panied by an exercise. ■ Trends in Ionization Energy General trends in first ionization energies of s- and p-block elements with position in the periodic table. Increase First ionization energyIn cr ea se ■ Factors Controlling Trends in Ionization Energies The ionization energy of an atom is always a balance between electron–nuclear attraction (which depends on Z) and electron–electron repulsion. KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. Afinidade eletrônica reason. In Groups 3A–5A, electrons are assigned to separate p orbitals (px, py, and pz). Beginning in Group 6A, however, two electrons are assigned to the same p or- bital. The fourth p electron shares an orbital with another electron and thus experi- ences greater repulsion than it would if it had been assigned to an orbital of its own: The greater repulsion experienced by the fourth 2p electron makes it easier to re- move, and each of the remaining p electrons has an orbital of its own. The usual trend resumes on going from oxygen to fluorine to neon, however, reflecting the in- crease in Z*. Active Figure 8.13 Electron Affinity Some atoms have an affinity, or “liking,” for electrons and can acquire one or more electrons to form a negative ion. The electron affinity, EA, of an atom is defined as the energy of a process in which an electron is acquired by the atom in the gas phase (Figure 8.14 and Appendix F). A(g)! e"(g) A"(g) E electron affinity, EA The greater the affinity an atom has for an electron, the more negative the value of EA will be. For example, the electron affinity of fluorine is "328 kJ/mol, a large value indicating an exothermic, product-favored reaction to form the anion, F". Boron has a much lower electron affinity for an electron, as indicated by a less neg- ative EA value of "26.7 kJ/mol. !¢S 2s 2p [Ne] 2s 2p [Ne] O (oxygen atom) !1314 kJ/mol O! (oxygen cation) ! e" 8.6 Atomic Properties and Periodic Trends 359 ■ Electron Affinity and Sign Conventions For a useful discussion of electron affinity, see J. C. Wheeler: “Electron affinities of the alkaline earth metals and the sign conven- tion for electron affinity.” Journal of Chemical Education, Vol. 74, pp. 123–127, 1997. Numerical values for EA are given in Appendix F. 0 "50 "100 "150 "200 "350 1 2 3 4 "250 "300 H El ec tr on a ff in it y (k J/ m ol ) 1A 2A 3A 4A 5A 6A 7A Group Period F O N C B Be Li Cl S P Si Al Mg Na Br Se As Ge Ga Ca K Active Figure 8.14 Electron affinity. The larger the affinity (EA) of an atom for an electron, the more negative the value. For numerical data, see Appendix F. (Data were taken from H. Hotop and W. C. Lineberger: “Binding ener- gies of atomic negative ions,” Journal of Physical Chemistry, Reference Data, Vol. 14, p. 731, 1985.) See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005. Raios iônicos Li!, 78 Li, 152 Na!, 98 Na, 186 K!, 133 K, 227 Rb!, 149 Rb, 248 Cs!, 165 Cs, 265 Be2!, 34 Be, 113 F", 133 F, 71 Cl", 181 Cl, 99 Br", 196 Br, 114 I", 220 I, 133 O2", 140 O, 66 N3", 146 N, 71 S2", 184 S, 104 Se2", 191 Se, 117 Te2", 211 Te, 143 Mg2!, 79 Mg, 160 Ca2!, 106 Ca, 197 Sr2!, 127 Sr, 215 Ba2!, 143 Ba, 217 Al3!, 57 Al, 143 Ga3!, 62 Ga, 122 In3!, 92 In, 163 Tl3!, 105 Tl, 170 1A 2A 3A 5A 6A 7A Main Group Metals Transition Metals Metalloids Nonmetals for Li. The loss of the 2s electron from Li leaves Li! with no electrons in the n 2 shell. The shrinkage will also be great when two or more electrons are removed, as for Al3! in which it exceeds 50%: You can also see by comparing Figures 8.11 and 8.15 that anions are always larger than the atoms from which they are derived. Here the argument is the opposite of that used to explain positive ion radii. The F atom, for example, has nine protons and nine electrons. On forming the anion, the nuclear charge is still !9, but now ten 3s 3p [Ne] 3s 3p [Ne] Al atom (radius # 143 pm) "3 electrons Al3! cation (radius # 57 pm) Li atom (radius # 152 pm) Li 1s 2s 1s 2s 78 pm Li! Li! cation (radius # 78 pm) " 1 electron 152 pm # 362 Chapter 8 Atomic Electron Configurations and Chemical Periodicity Active Figure 8.15 Relative sizes of some common ions compared with neutral atom size. Radii are given in picometers (1 pm# 1$ 10"12 m). (Data taken from J. Emsley: The Elements, 3rd ed., Oxford, Clarendon Press, 1998.) See the General ChemistryNow CD-ROM or website to explore an interactive version of this figure accompanied by an exercise. KOTZ, J. C.; TREICHEL, P. M. Química e reações químicas. 2005.
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