Chapter

1
Chapter 8 - Electron Configuration and
Chemical Periodicity
8.1 Development of the Periodic Table 8.2
Characteristics of Many-Electron Atoms 8.3 The
Quantum-Mechanical Model and the Periodic
Table 8.4 Trends in Some Key Periodic Atomic
Properties 8.5 The Connection Between Atomic
Structure and Chemical Reactivity
2
Mendeleevs Predicted vs Actual Properties
of Element 32 - Germanium
Property Predicted
Properties Actual Properties
Atomic Mass 72
72.59
Appearance Gray Metal
Gray Metal
Density 5.5
g/cm3 5.35 g/cm3
Molar volume 13 cm3 /mol
13.22 cm3/mol
Specific heat capacity 0.31 J/g K
0.32 J/g K
Oxide density 4.7 g/cm3
4.23 g/cm3
Sulfide formula and ES2 insoluble
in GeS2 insoluble in
solubility H2O soluble in
H2O soluble in
aqueous (NH4)S
aqueous (NH4)S
Chloride formula ECl4
GeCl4
(boiling point) lt 100oC
84oC
Chloride density 1.9 g/cm3
1.844 g/cm3
Element preparation reduction of
K2EF6 Reduction of
with sodium
K2GeF6 with sodium
3
Observing the Effect of Electron Spin
Fig. 8.1
4
Summary of Quantum Numbers of Electrons in Atoms
Name Symbol Permitted Values
Property
Principal n Positive integers
(1,2,3, etc.) Orbital energy

(size) Angular
l Integers from 0 to n - 1
Orbital shape (the l momentum

values 0, 1, 2, and 3

correspond to the s,

p, d, and f orbitals) Magnetic
ml Integers from -l to 0 to l
Orbital orientation Spin ms
1/2 or -1/2
Direction of e- spin


Table 8.2
5
Quantum Numbers - I

• 1) Principal Quantum Number n
• Also called the energy quantum number,
  indicates the approximate distance from
  the nucleus .
• Denotes the electron energy shells around the
  atom, and is derived directly from the
  Schrodinger equation.
• The higher the value of n , the greater the
  Energy of the orbital, and hence the energy of
  electrons in that orbital.
• Positive integer values of n 1 , 2 , 3 , etc.

6
Quantum Numbers - II

• 2) Azimuthal
• Denotes the different energy sublevels within the
  main level n
• Also indicates the shape of the orbitals around
  the nucleus.
• Positive interger values of L are 0
  ( n-1 )
• n 1 , L 0 n 2 ,
  L 0 and 1
• n 3 , L 0 , 1 , 2

7
Quantum Numbers - III

• 3) Magnetic Quantum Number - mL Also called
  the orbital orientation Quantum
• denotes the direction or orientation in a
  magnetic field - Or it denotes the different
  magnetic geometriesound the nucleus - three
  dimensional space
• values can be positive and negative (-L 0
  L)
• L 0 , mL 0 L 1 , mL
  -1,0,1
• L 2 , mL -2,-1,0,1,2

8
Quantum Numbers - IV

• 4) Spin Quantum Number - ms - gives the spin of
  the electron or -
• The values of the spin are either
• 1 / 2 or - 1 / 2
• n 1 L 0 mL 0 ms 1/ 2 and -
  1/ 2
• n 2 L 0 mL 0 ms 1/ 2 and -
  1/ 2
• L 1 mL -1 ms 1/ 2
  and - 1/ 2
• mL 0 ms 1/ 2
  and - 1/ 2
• mL 1 ms
  1/ 2 and - 1/ 2

9
Spectral Evidence of Energy-Level Splitting
in Many-Electron Atoms
Fig. 8.2
10
Fig. 8.4
11
Pauli Exclusion Principle

• Each electron in an atom must have a unique set
  of quantum numbers !
• Only two electrons can be described by the same
  orbital and these two electrons must have
  opposite spin.

12
As a Result of the Pauli Exclusion Principle

• Electrons with the same spin keep apart in space
  whereas electrons of opposite spin may occupy the
  same region of space.

13
Quantum Numbers - V

• n 1 L 0 mL 0 ms 1/ 2
  - 1/ 2
• n 2 L 0 mL 0 for
  all orbitals
• L 1 mL -1 , 0 , 1
• n 3 L 0 mL 0
• L 1 mL -1 , 0 , 1
• L 2 mL - 2 , -1 , 0 , 1
  , 2
• n 4 L 0 mL 0
• L 1 mL -1 , 0 1
• L 2 mL - 2 , -1 , 0 , 1
  , 2
• L 3 mL - 3 , - 2 , - 1 ,
  0, 1,2 ,3

14
Quantum Numbers - VI
Allowed Values
n
1 2 3
4
L
0 0 1 0 1
2 0 1 2 3
mL
0 0 -1 0 1 0 -1 0 1
0 -1 0 1
-2 -1
0 1 2 -2 -1 0 1 2
-3 -2
-1 0 1 2 3
ms
All or - 1/2 spin
1/2 -1/2
15
Quantum Numbers - VII Noble Gases
Electron Orbitals
Number of Electrons Element
1s2
2
He
1s2 2s22p6
10
Ne
1s2 2s22p6 3s23p6
18 Ar
1s2 2s22p6 3s23p6 4s23d104p6
36 Kr
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
54 Xe
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
6s24f14 5d106p6 86 Rn
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6
6s24f145d106p6 118 ?

7s25f146d107p6
16
Electron Configuration of Helium and Lithium

• He 1s2
• n 1 L 0 mL 0 ms
  1/ 2
• n 1 L 0 mL 0 ms -
  1/ 2
• Li 1s2 2s1
• n 1 L 0 mL 0 ms
  1/ 2
• n 1 L 0 mL 0 ms -
  1/ 2
• n 2 L 0 mL 0 ms -
  1/ 2

17
Orbital Box Diagrams - I H Be
Element Symbol Electron
Orbital Box Diagrams
Configuration
Hydrogen H 1s1
Helium He 1s2 Lithium
Li 1s22s1 Beryllium
Be 1s22s2
1s
2s
1s
2s
2s
1s
1s
2s
18
Fig. 8.5
19
Hunds Rule

• For an atom in its ground-state configuration,
  all unpaired electrons have the same spin
  orientation.
• Therefore electrons tend to occupy all free
  orbitals and not pair up, so that their spins all
  add up to produce a general vector for the atom.

20
Orbital Occupancy for the First 10 Elements, H
through Ne
Fig. 8.6
21
Orbital Box Diagrams - II B Ne
B (5 e-) 1s2 2s2 2p1
C (6 e-) 1s2 2s2 2p2
N (7 e-) 1s2 2s2 2p3
O (8 e-) 1s2 2s2 2p4
F (9 e-) 1s2 2s2 2p5
Ne (10 e-) 1s2 2s2 2p6
22
Valence and Core Electrons

• Valence Electrons - Those electrons outside of a
  closed electron shell. These electrons take part
  in chemical reactions.
• Core Electrons - The electrons in the closed
  shells. They cannot take part in chemical
  reactions.
• Sodium 11 electrons
• Valence electrons Ne 3s 1 --- one
• Core electrons 1s 2 2s 2 2p 6 ---
  Ten
• Chlorine 17 electrons
• Valence electrons Ne 3s 2 3p 5----
  seven
• Core 2 2s 2 2p 6 ---- Ten

23
Quantum Numbers and the Number of Electrons

• n L m s
  e-
• 
  
• 1 0 0 (1s) 1/2 -
  1/2 2 2
• 2 0 0 (2s) 1/2 -1/2
  2 4
• 1 -1,0,1 (2p) 1/2-1/2
  6 10
• 3 0 0 (3s) 1/2-1/2
  2 12
• 1 -1,0,1 (3p) 1/2-1/2
  6 18
• 2 -2,-1,0,1,2(3d) 1/2-1/2
  10 28
• 4 0 0 (4s) 1/2-1/2
  2 30
• 1 -1,0,1 (3p) 1/2-1/2
  6 36
• Denotes a noble gas !!!

24
Order of Electron Filling
25
Electron Configuration - I

• H 1s 1
• He 1s 2
  He
• Li 1s2 2s 1
  He 2s 1
• Be 1s2 2s 2 He 2s 2
• B 1s2 2s 2 2p 1 He 2s 2 2p
  1
• C 1s 2 2s 2 2p 2 He 2s 2 2p
  2
• N 1s 2 2s 2 2p 3 He 2s 2 2p
  3
• O 1s 2 2s 2 2p 4 He 2s 2 2p
  4
• F 1s 2 2s 2 2p 5 He 2s 2
  2p 5
• Ne 1s 2 2s 2 2p 6 He 2s 2 2p6
  Ne

26
Electron Configuration - II

• Na Ne 3s 1
• Mg Ne 3s 2
• Al Ne 3s 2 3p 1
• Si Ne 3s 2 3p 2
• P Ne 3s 2 3p 3
• S Ne 3s 2 3p 4
• Cl Ne 3s 2 3p 5
• Ar Ne 3s 2 3p6 Ar

27
Condensed Ground-State Electron
Configurations in the First Three Periods
Fig. 8.7
28
Orbital Box Diagrams - III Na
Ar
Atomic Number Orbital Box
Condensed Electron Element
Diagrams(3s3p)
Configuration
11 Na
He 3s1 12
Mg
He 3s2 13 Al

He 3s23p1 14 Si

He 3s23p2 15 P

He 3s23p3 16 S

He 3s23p4 17 Cl

He 3s23p5 18 Ar

He 3s23p6
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3px
3py
3pz
3s
3py
3px
3pz
3s
3py
3px
3pz
3s
3py
3px
3pz
3py
3px
3pz
3s
29
Fig. 8.8
30
Electron Configuration - III

• K Ar 4s 1
• Ca Ar 4s 2 Or this order
  is OK !
• Sc Ar 4s 2 3d 1
  Ar 3d 1 4s 2
• Ti Ar 4s 2 3d 2
  Ar 3d 2 4s 2
• V Ar 4s 2 3d 3
  Ar 3d 3 4s 2
• Cr Ar 4s 1 3d 5
• Mn Ar 4s 2 3d 5
• Fe Ar 4s 2 3d 6 Either order will
  be OK !
• Co Ar 4s 2 3d 7 But its normally
  best to
• Ni Ar 4s 2 3d 8 put the one filling
  last!!!
• Cu Ar 4s 1 3d 10
• Zn Ar 4s 2 3d 10

Anomalies to Filling
Anomalies to Filling
31
Orbital Box Diagram - IV Sc Zn
4s 3d
Z 21 Sc Ar 4s2 3d1
Z 22 Ti Ar 4s 2 3d 2
Z 23 V Ar 4s 2 3d 3
Z 24 Cr Ar 4s1 3d 5
Z 25 Mn Ar 4s 2 3d 5
Z 26 Fe Ar 4s 2 3d 6
Z 27 Co Ar 4s 2 3d 7
Z 28 Ni Ar 4s 2 3d 8
Z 29 Cu Ar 4s 1 3d 10
Z 30 Zn Ar 4s 2 3d 10
32
Electron Configuration - IV

• Ga Ar 4s 2 3d 10 4p 1
• Ge Ar 4s 2 3d 10 4p 2
• As Ar 4s 2 3d 10 4p 3
• Se Ar 4s 2 3d 10 4p 4
• Br Ar 4s 2 3d 10 4p 5
• Kr Ar 4s 2 3d 10 4p 6 Kr

33
Electron Configuration - V

• Rb Kr 5s 1
• Sr Kr 5s 2
• Y Kr 5s 24d 1
• Zr Kr 5s 2 4d 2
• Nb Kr 5s 1 4d 4
• Mo Kr 5s 1 4d 5
• Tc Kr 5s 2 4d 6
• Ru Kr 5s 1 4d7
• Rh Kr 5s 1 4d 8
• Pd Kr 4d 10
• Ag Kr 5s 1 4d 10
• Cd Kr 5s 2 4d 10

Anomalies to Filling
34
Electron Configuration - VI

• In Kr 5s 2 4d 10 5p 1
• Sn Kr 5s 2 4d 10 5p 2
• Sb Kr 5s 2 4d 10 5p 3
• Te Kr 5s 2 4d 10 5p 4
• I Kr 5s 2 4d 10 5p 5
• Xe Kr 5s 2 4d 10 5p 6 Xe

35
Electron Configuration - VII

• Cs Xe 6s 1
• Ba Xe 6s 2
• La Xe 6s2 5d 1
• Ce Xe 6s 2 5d 1 4f 1
• Pr Xe 6s 2 4f 3
• Nd Xe 6s 2 4f 4
• Pm Xe 6s 2 4f 5
• Sm Xe 6s 2 4f 6
• Eu Xe 6s 2 4f 7
• Gd Xe 6s 2 3d 1 4f 7
• Tb Xe 6s 2 4f 9
• Dy Xe 6s 2 4f 10
• Ho Xe 6s 2 4f 11
• Er Xe 6s 2 4f 12
• Tm Xe 6s 2 4f 13
• Yb Xe 6s 2 4f 14
• Lu xe 6s 2 3d 1 4f 14

Anomalies to Filling
36
Electron Configuration - VIII

• Hf Xe 6s 2 4f 14 5d 2
• Ta Xe 6s 2 4f 14 5d 3
• W Xe 6s 2 4f 14 5d 4
• Re Xe 6s 2 4f 14 5d 5
• Os Xe 6s 2 4f 14 5d 6
• Ir Xe 6s 2 4f 14 5d 7
• Pt Xe 6s 1 4f 14 5d 9
  
• Au Xe 6s 1 4f 14 5d 10
  
• Hg Xe 6s 2 4f 14 5d 10
• Tl Xe 6s 2 4f 14 5d 10 6p 1
• Pb Xe 6s 2 4f 14 5d 10 6p 2
• Bi Xe 6s 2 4f 14 5d 10 6p 3
• Po Xe 6s 2 4f 14 5d 10 6p 4
• At Xe 6s 2 4f 14 5d 10 6p 5
• Rn xe 6s 2 4f 14 5d 10 6p 6 Rn

Anomalies to Filling
37
Electron Configuration - IX

• Fr Rn 7s 1
• Ra Rn 7s 2
• Ac Rn 7s 2 6d 1
• Th Rn 7s 2 6d 2
• Pa Rn 7s 2 5f 2 6d 1
• U Rn 7s 2 5f 3 6d 1
• Np Rn 7s 2 5f 4 6d 1
• Pu Rn 7s 2 5f 6
• Am Rn 7s 2 5f 7
• Cm Rn 7s 2 5f 7 6d 1
• Bk Rn 7s 2 5f 9
• Cf Rn 7s 2 5f 10
• Es Rn 7s 2 5f 11
• Fm Rn 7s 2 5f 12
• Md Rn 7s 2 5f 13
• No Rn 7s 2 5f 14
• Lr Rn 7s 2 5f 14 6d 1

Anomalies to Filling
38
The Periodic Table of the Elements
Electronic Structure
He
H
Ne
F
O
N
C
B
Li
Be
Ar
Cl
S
P
Si
Al
Na
Mg
Kr
Zn
Cu
Ni
Co
Fe
Mn
Cr
V
Ti
Sc
Br
Se
As
Ge
Ga
K
Ca
Xe
Cd
Ag
Pd
Rh
Ru
Tc
Mo
Nb
Zr
Y
I
Te
Sb
Sn
In
Rb
Sr
Rn
Hg
Au
Pt
Ir
Os
Re
W
Ta
Hf
La
At
Po
Bi
Pb
Tl
Cs
Ba
Ac
Rf
Ha
Fr
Ra
Sg
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
S Orbitals
P Orbitals f Orbitals
d Orbitals
39
The Periodic Table of the Elements
Anomolies to Electron Filling
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Ar
Al
Si
P
S
Cl
K
Ca
Sc
Ti
V
Mn
Fe
Co
Ni
Zn
Ga
Ge
As
Se
Br
Kr
Cu
Cr
Rb
Sr
Y
Zr
Tc
Cd
Xe
I
Te
Sb
Sn
In
Nd
Mo
Ru
Rh
Pd
Ag
Cs
Ba
Hf
Ta
W
Re
Os
Ir
Hg
Rn
At
Po
Bi
Pb
Tl
La
Pt
Au
Fr
Ra
Rf
Sg
Ac
Du
Bo
Ha
Me
Pr
Nd
Pm
Sm
Eu
Tb
Dy
Ho
Er
Tm
Yb
Lu
Ce
Gd
Pu
Am
Bk Cf
Es
Fm
Md
No
Lr
Th
Pa
U
Np
Cm
Anomalous Electron Filling
40
A Periodic Table ofPartial Ground-State Electron
Configurations
Fig. 8.9
41
Fig. 8.10
42
Electronic Configuration Ions

• Na 1s 2 2s 2 2p 6 3s 1 Na 1s 2
  2s 2 2p 6
• Mg 1s 2 2s 2 2p 6 3s 2 Mg2 1s 2 2s
  2 2p6
• Al 1s 2 2s 2 2p 6 3s 2 3p 1 Al3 1s 2
  2s 2 2p 6
• O 1s 2 2s 2 2p 4 O- 2 1s
  2 2s 2 2p 6
• F 1s 2 2s 2 2p 5 F- 1
  1s 2 2s 2 2p 6
• N 1s 2 2s 2 2p 3
  N- 3 1s 2 2s 2 2p 6

43
Fig. 8.11
44
Atomic Radii of the Main-Group and Transition
Elements
Fig. 8.12
45
Fig. 8.13
46
Ranking Elements by Size
Problem Rank the following elements in each
group according to decreasing
size ( largest first!) a) Na, K, Rb b)
Sr, In, Rb c) Cl, Ar, K d) Sr, Ca,
Rb Plan Find their relative position in the
periodic table and apply trends! Solution
a) Rb gt K gt Na These elements are all
alkali metals and the
elements increase in size as you go down the
group. b) Rb gt Sr gt In These elements are
in Period 5 and the size
decreases as you go across the
period. c) K gt Cl gt Ar These elements
border a noble gas, and it is the
smallest
diameter. d) Rb gt Sr gt Ca These elements are
near each other, Sr is beneath
Ca therefore it is larger and Rb is next to
Sr and larger.
47
Periodicity of First Ionization Energy (IE1)
Fig. 8.14
48
Fig. 8.15
49
Fig. 8.16
50
Successive Ionization Energies
Valence Z Element
Electrons IE1 IE2 IE3 IE4 IE5
IE6 IE7
3 Li 1 0.52 7.30
11.81 4 Be 2
0.92 1.76 14.85 21.01 5 B
3 0.80 2.43 3.66 25.02
32.82 6 C 4
1.09 2.35 4.62 6.22 37.83 47.28 7
N 5 1.40 2.86 4.58
7.48 9.44 53.27 64.36 8 O
6 1.31 3.39 5.30 7.47
10.98 13.33 71.33 9 F
7 1.68 3.37 6.05 8.41 11.02
15.16 17.87
51
Ranking Elements by First Ionization Energy
Problem Using the Periodic table only, rank the
following elements in each
of the following sets in order of increasing
IE! a) Ar, Ne, Rn b) At, Bi, Po c)
Be, Na, Mg d) Cl, K, Ar Plan Find their
relative positions in the periodic table and
apply trends! Solution
a) Rn, Ar,Ne These elements are all noble
gases and their IE
decreases as you go down the group. b)
Bi, Po, At These elements are all Period 6
elements and the IE
increases from the left to the
right. c) Na, Mg, Be These elements are close
to each other, Be Mg
are in the same group, Be is higher than Mg
Na
is next to Mg lower in IE. d) K, Cl, Ar These
elements bracket the noble gas Ar, and Cl would
be lower than Ar and
K would be lower yet!
52
Identifying Elements by Its Successive
Ionization Energies
Problem Given the following series of
ionization energies (in kJ/mol)
for an element in period 3, name the element and
write its electron
configuration
IE1 IE2 IE3
IE4 580
1,815 2,740
11,600 Plan Examine the values to find the
largest jump in ionization energy,
which occurs after all valence electrons have
been removed. Use the periodic
table! Solution
The largest jump in IE occurs after IE3 so the
element has 3 valence electrons thus it is
Aluminum ( Al, Z13), its electron configuration
is
1s2 2s2 2p6 3s2 3p1
53
Electron Affinities of the Main-Group Elements
Fig. 8.17
54
Trends in Three Atomic Properties
Fig 8.18
55
Isoelectronic Atoms and Ions

• H- 1 He Li Be2
• N- 3 O- 2 F- Ne Na Mg2
  Al3
• P- 3 S- 2 Cl- Ar K Ca2
  Sc3 Ti4
• As- 3 Se- 2 Br- Kr Rb Sr2
  Y3 Zr4
• Sb- 3 Te- 2 I- Xe Cs Ba2
  La3 Hf4

56
Trends in Metallic Behavior
Fig. 8.19
57
Fig. 8.20
58
Fig. 8.21
59
The Trends in Acid-Base Behavior of Elemental
Oxides
Fig. 8.22
60
Fig. 8.23
61
Electron Configurations of Ions of Main Group
Elements - and Their Charge
Problem Write reactions with condensed electron
configurations to show the formation of the
common ions of the following elements a)
Sulfur (Z16) b) Barium (Z56) c)
Antimony (Z 51) Plan We identify the elements
position in the periodic table, and keep
two generalizations in mind Ions of
elements in groups 1A,2A,6A,and 7A are typically
isoelectronic with the nearest noble gas.
Metals in groups 3Ato 5A can lose their ns
or their ns and np electrons. Solution
.
.
a) S Ne 3s23p4 2 e -
S2- Ne 3s23p6 (same as Ar) b) Ba (Xe 6s2)

Ba2 Xe 2 e - c) Sb Kr 4d105s25p3
Sb3 Kr 4d105s2 3 e
- Sb Kr 4d105s25p3
Sb5 Kr 4d10 5 e -
62
Fig. 8.24
63
Pseudo - Noble Gas Electron Configurations
Elements in groups 3A, 4A, and 5A can form
cations by losing enough electrons to leave a
pseudo noble gas configuration. By losing
electrons and leaving a filled d orbital, which
is quite stable!
Sn Kr 5s24d105p2
Sn4 Kr 4d10 4 e - Sn Kr 5s24d105p2
Sn2 Kr 5s24d10
2 e - Pb Xe 4f145d106s26p2
Pb2 Xe 4f145d106s2 2 e- Pb Xe
4f145d106s26p2 Pb4 Xe
4f145d10 4 e- As Ar 3d104s24p3
As3 Ar 3d104s2 3
e- As Ar 3d104s24p3
As5 Ar 3d10 5 e- Sb Kr 4d105s25p3
Sb3 Kr
4d105s2 3 e- Sb Kr 4d105s25p3
Sb5 Kr 4d10 5 e-
64
Magnetic Properties

• Paramagnetic - An atom or ion which has unpaired
  electrons, which add up to give a spin vector.
  They are thereby attracted by a magnetic field.
• Diamagnetic - An atom or ion with all electrons
  paired and with no net spin.

65
Apparatus for Measuring the Magnetic Behavior of
a Sample
Fig. 8.25
66
Examples of Elements and Ions That are
Paramagnetic
Ti Ar4s23d2 Ti2
Ar 3d2 2 e -
a)
4s 3d
4s 3d
c)
Cu Ar 4s1 3d10 Cu1
Ar 3d10 1 e -
Cu or Zn2
4s 3d
Zn Ar 4s2 3d10 Zn2
Ar 3d10 2 e -
67
Fig. 8.26
68
Fig. 8.27
69
Ranking Ions According to Size
Problem Rank each set of Ions in order of
increasing size. a) K, Rb, Na b) Na,
O2-, F - c) Fe2, Fe3 Plan We find the
position of each element in the periodic table
and apply the ideas of size i) size increases
down a group, ii) size decreases across a period
but increases from cation to anion. iii) size
decreases with increasing positive (or
decreasing negative) charge in an isoelectronic
series. iv) cations of the same element decreases
in size as the charge increases. Solution
70
Ranking Ions According to Size
Problem Rank each set of Ions in order of
increasing size. a) K, Rb, Na b) Na,
O2-, F - c) Fe2, Fe3 Plan We find the
position of each element in the periodic table
and apply the ideas of size i) size increases
down a group, ii) size decreases across a period
but increases from cation to anion. iii) size
decreases with increasing positive (or
decreasing negative) charge in an isoelectronic
series. iv) cations of the same element decreases
in size as the charge increases. Solution a)
since K, Rb, and Na are from the same group
(1A), they increase in size down the
group Na lt K lt Rb
71
Ranking Ions According to Size
Problem Rank each set of Ions in order of
increasing size. a) K, Rb, Na b) Na,
O2-, F - c) Fe2, Fe3 Plan We find the
position of each element in the periodic table
and apply the ideas of size i) size increases
down a group, ii) size decreases across a period
but increases from cation to anion. iii) size
decreases with increasing positive (or
decreasing negative) charge in an isoelectronic
series. iv) cations of the same element decreases
in size as the charge increases. Solution a)
since K, Rb, and Na are from the same group
(1A), they increase in size down the
group Na lt K lt Rb b) the ions Na, O2-,
and F- are isoelectronic. O2- has lower Zeff than
F-, so it is larger. Na is a cation, and
has the highest Zeff, so it is smaller
Na lt F- lt O2-
72
Ranking Ions According to Size
Problem Rank each set of Ions in order of
increasing size. a) K, Rb, Na b) Na,
O2-, F - c) Fe2, Fe3 Plan We find the
position of each element in the periodic table
and apply the ideas of size i) size increases
down a group, ii) size decreases across a period
but increases from cation to anion. iii) size
decreases with increasing positive (or
decreasing negative) charge in an isoelectronic
series. iv) cations of the same element decreases
in size as the charge increases. Solution a)
since K, Rb, and Na are from the same group
(1A), they increase in size down the
group Na lt K lt Rb b) the ions Na, O2-,
and F- are isoelectronic. O2- has lower Zeff than
F-, so it is larger. Na is a cation, and
has the highest Zeff, so it is smaller
Na lt F- lt O2- c) Fe2 has a lower charge than
Fe3, so it is larger Fe3 lt Fe2