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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.