8-1

Figure 8.6 Orbital Filling Order

Chapter 8 Electron Configuration &Chemical Periodicity One additional quantum number is needed to describe a property of an electron in an atomic orbital.

4. Spin Quantum Number, ms

ms = + 12 or – 12

An electron in an orbital is described by its set of FOUR quantum numbers.

“shell”: orbitals with the same n value “subshell”: orbitals with the same n & l values

e.g., the n=3 shell contains the 3s subshell, 3p subshell, and 3d subshell.

Pauli Exclusion Principle No two electrons in an atom can have the same set of four quantum numbers.

Therefore: a maximum of TWO electrons per orbital. (A third electron in an orbital would have to have the same set of four quantum numbers as one of the two electrons already there.) Hund’s Rule

When orbitals of equal energy (“degenerate”) are available, the electron configuration of lowest energy has the maximum number of electrons with the same spin.

8-2

General Principles of Electron Configuration Note similar electron configurations within a group e.g., Group 5A: ns2np3

Orbital Filling Order & Periodic Table The periodic table is our best guide to filling orbitals and writing electron configurations!

1A (1)

8A (18)

1 H

1.008

2A (2)

3A (13)

4A (14)

5A (15)

6A (16)

7A (17)

2 He 4.003

3 Li

6.941

4 Be 9.012

5 B

10.81

6 C

12.01

7 N

14.01

8 O

16.00

9 F

19.00

10 Ne 20.18

11 Na 22.99

12 Mg 24.31

3B (3)

4B (4)

5B (5)

6B (6)

7B (7)

(8)

8B (9)

(10)

1B (11)

2B (12)

13 Al

26.98

14 Si

28.09

15 P

30.97

16 S

32.07

17 Cl

35.45

18 Ar 39.95

19 K

39.10

20 Ca 40.08

21 Sc

44.96

22 Ti

47.88

23 V

50.94

24 Cr 52.00

25 Mn 54.94

26 Fe 55.85

27 Co 58.93

28 Ni

58.69

29 Cu 63.55

30 Zn 65.39

31 Ga 69.72

32 Ge 72.61

33 As 74.92

34 Se

78.96

35 Br 79.90

36 Kr 83.80

37 Rb 85.47

38 Sr

87.62

39 Y

88.91

40 Zr 91.22

41 Nb 92.91

42 Mo 95.94

43 Tc (98)

44 Ru

101.07

45 Rh

102.91

46 Pd

106.42

47 Ag

107.87

48 Cd

112.41

49 In

114.82

50 Sn

118.71

51 Sb

121.76

52 Te

127.60

53 I

126.90

54 Xe

131.29 55 Cs

132.91

56 Ba

137.33

71 Lu

174.97

72 Hf

178.49

73 Ta

180.95

74 W

183.84

75 Re

186.21

76 Os

190.23

77 Ir

192.22

78 Pt

195.08

79 Au

197.00

80 Hg

200.59

81 Tl

204.38

82 Pb 207.2

83 Bi

208.98

84 Po (209)

85 At (210)

86 Rn (222)

87 Fr (223)

88 Ra [226]

103 Lr (262)

104 Rf (267)

105 Db (268)

106 Sg (271)

107 Bh (270)

108 Hs (269)

109 Mt (278)

110 Ds (281)

111 Rg (281)

112 Cn (285)

113 Nh (286)

114 Fl

(289)

115 Mc (289)

116 Lv (293)

117 Ts (294)

118 Og (294)

57 La

138.91

58 Ce

140.12

59 Pr

140.91

60 Nd

144.24

61 Pm (145)

62 Sm 150.36

63 Eu

151.96

64 Gd 157.25

65 Tb

158.93

66 Dy

162.50

67 Ho

164.93

68 Er

167.26

69 Tm 168.93

70 Yb

173.04 89 Ac [227]

90 Th

232.04

91 Pa

231.04

92 U

238.03

93 Np [237]

94 Pu (244)

95 Am (243)

96 Cm (247)

97 Bk (247)

98 Cf (251)

99 Es (252)

100 Fm (257)

101 Md (258)

102 No (259)

Inner (core) electrons: those in the previous noble gas configuration plus those in filled d and f subshells (i.e., d10 and f 14).

Valence electrons: electrons in highest energy levels (largest n value). o For A-group elements, the valence electrons are the outer electrons. o For B-group elements (TMs), some d electrons are sometimes counted as valence electrons.

PROTIP You may notice that La in this table appears in the f block. Doing so gives correct predictions of configurations for most of the other f block elements.

Figure 8.10

8-3

Periodic Trends Atomic Radius

Usually measured using distances between nuclei of bonded atoms (e.g., “bonding atomic radius”); see Figure 8.14. Trends Going down a group, atomic radius increases. This is the result of the increasing principal quantum number of the valence electrons (larger n larger size of orbital). Going across a period (left to right), atomic radius decreases. This is the result of an increasing nuclear charge (Z) attracting electrons with the same principal quantum number.

Figure 8.13: Atomic radii in picometers.

Ionization Energy

Ionization energy is defined as the minimum energy required to remove 1 mol of electrons from 1 mol of gaseous atoms/ions. First ionization energy, IE1, is defined as the amount of

energy required to remove 1 mol of electrons from gaseous neutral atoms. X(g) →X+(g) + e–

Second ionization energy, IE2, is defined as the amount of energy required to remove 1 mol of electrons from gaseous ions with a +1 charge, producing ions with a 2+ charge. X+(g) →X2+(g) + e–

Trends Going down a group, first ionization energy decreases, as the result of the increasing principal quantum number of the valence electrons (larger n larger size of orbital). Going across a period (left to right), first ionization energy increases. This is the result of an increasing nuclear charge (Z) attracting electrons with the same principal quantum number. The greater attraction for the electrons makes them harder to remove.

Figure 8.16: First ionization energies in kJ∙mol–1. For each element, IE1 < IE2 < IE3 and so forth. There is a sharp jump in IE when an inner-shell electron is removed. For example, removing the four valence electrons from C requires progressively more energy for each electron. Removing the fifth electron from C requires MUCH more energy because at that point, you are removing core electrons.

Table 8.5

8-4

Electron Affinity

Electron affinity is defined as the energy change with the addition of 1 mole of electrons to 1 mole of gaseous atoms/ions. X(g) + e– →X–(g)

Gross Trends There are many exceptions to these gross trends, which result from the arrangement of electrons in orbitals and the special stability of half-filled orbitals (Hund’s rule), among other things. Going down a group, electron affinity generally decreases. Going across a period (left to right), electron affinity generally increases. For the noble gases, however, the EA is negative because addition of an electron would disrupt the already-stable arrangement of electrons. The additional electron would have to reside in a higher-energy subshell.

Figure 8.18, Electron affinities in kJ∙mol–1.

Metallic Behavior

Metals: typically shiny solids, moderate to high melting points, good thermal/electrical conductors, tend to lose electrons in reactions with nonmetals

Nonmetals: everything the opposite of metals Metalloids: share some properties of metals and nonmetals.

Nonmetal Oxides (some solids, some liquids, some gases) nonmetal oxide + water ACID e.g., SO3 + H2O H2SO4

Main Group Metal Oxides (always solids) metal oxide + water metal hydroxide (a BASE) e.g., K2O + H2O 2KOH

Transition metal oxides may be acidic, basic, or either! High oxidation number results in more covalent bonding. e.g., CrVIO3 + H2O H2CrO4 “chromic acid”, a strong acid—note similarity to SO3 which also has a +6 O.N.

SUMMARY of Periodic Trends