Electronic configurations and the Periodic Table

1
Electronic Configurations and the Periodic Table
5.1 Relative Energies of Orbitals 5.2 Electronic
Configurations of Elements 5.3 The Periodic
Table 5.4 Ionization Enthalpies of
Elements 5.5 Variation of Successive Ionization
Ethalpies with Atomic Numbers 5.4 Atomic
Size of Elements
2
Relative Energies of Orbitals
3
In one-electron systems (e.g. H, He), there are
no interactions(no shielding effects) between
electrons. All subshells(s, p, d, f,) of the
same principal quantum shell have the same
energy. The subshells are said to be degenerate.
4
Evidence
In the Lyman series, only one spectral line is
observed for the transition from n 2 to n 1.
2s and 2p subshells are degenerate
5
In multi-electron systems, there are
interactions(shielding effects) between
electrons. Different subshells of the same
principal quantum shell occupy different energy
levels. The energies of subshells or orbitals
follow the order s lt p lt d lt f
6
Relative energies of orbitals
5.1 Relative energies of orbitals (SB p.106)
7
Relative energies of orbitals
5.1 Relative energies of orbitals (SB p.106)
8
Relative energies of orbitals
5.1 Relative energies of orbitals (SB p.106)
9
Relative energies of orbitals
5.1 Relative energies of orbitals (SB p.106)
10
Both 4s and 3d electrons are shielded from the
nuclear attraction by the inner core (2,8,8)
11
4s electron is more penetrating than 3d electron,
spending more time closer to the nucleus.
4s electron experiences stronger nuclear
attraction
4s electron is more stable.
12
Three rules to build up electronic
configurations 1. Aufbau (building up)
Principle 2. Hunds Rule 3. Paulis Exclusion
Principle
13

1. Aufbau (building up) Principle

Electrons enter the orbitals in order of
ascending energy.
14
Numbers read downwards
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
15
1s,
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
16
1s, 2s,
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
17
1s, 2s, 2p, 3s,
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
18
1s, 2s, 2p, 3s, 3p, 4s,
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
19
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
20
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
21
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s,
4f, 5d, 6p, 7s
s p d f g h i
1 2 3 4 5 6 7
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
22
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s,
4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s
s p d f g h i
1 2 3 4 5 6 7 8
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7
7
23
Building up of electronic configurations
5.1 Relative energies of orbitals (SB p.106)
24
2. Hunds rule - Orbitals of the same energy
must be occupied singly and with the same spin
before pairing up of electrons occurs.
Electrons-in boxes diagram
25
3. Paulis exclusion principle - Electrons
occupying the same orbital must have
opposite spins.
26
Electronic Configurations of Elements
27
Ways to Express Electronic Configurations
5.2 Electronic configurations of elements (SB
p.108)
1. The s, p, d, f notation
Na 1s2, 2s2, 2p6, 3s1 1s2, 2s2, 2px2, 2py2,
2pz2, 3s1
28
Q.17(a)
K 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Q.17(b)
Fe 1s2, 2s2, 2p6, 3s2, 3p6, 3d6, 4s2
29
2. Using a noble gas core
Na Ne 3s1 Ca Ar 4s2
30
Q.18(a)
Si Ne 3s2, 3p2
Q.18(b)
V Ar 3d3, 4s2
31
3. Electrons in Boxes representation

1. All boxes should be labelled
2. Boxes of the same energies are put together.

32
Q.19
33
Q.20(a)
Phosphorus
34
Q.20(b)
Chromium
The half-filled 3d subshell has extra stability
due to the more symmetrical distribution of
charge. The energy needed to promote an electron
from 4s to 3d is more than compensated by the
energy released from the formation of half-filled
3d subshells.
35
Q.20(b)
Ar 3d4, 4s2 ? Ar 3d5, 4s1 energy
Chromium
36
Q.20(c)
Copper
The full-filled 3d subshell has extra stability
due to the more symmetrical distribution of
charge. The energy needed to promote an electron
from 4s to 3d is more than compensated by the
energy released from the formation of full-filled
3d subshells.
37
Ar 3d9, 4s2 ? Ar 3d10, 4s1 energy
Q.20(c)
Copper
38
Silicon
Ne
Empty orbital(s) in a partially filled subshell
should be shown
39
Silicon
energy
?
Energy difference 3p 3s gt 3d 4s
40
21(a)
21(b) Ar
41
S2? , Cl? , Ar , K , Ca2
Same electronic configurations Isoelectronic
Q.22
42
Represented by electrons-in-boxes diagrams
5.2 Electronic configurations of elements (SB
p.110)
43
5.2 Electronic configurations of elements (SB
p.110)
44
Building up of electronic configurations
http//www.chemcollective.org/applets/pertable.php
45
A brief history of the Periodic Table
Ancient Greece, Aristotle - Four elements Fire,
Water, Air, Earth,
46
A brief history of the Periodic Table
Ancient Greece, Aristotle - Four elements Air,
Fire, Earth, Water
Buddha ???????
??
Quintessence (The fifth element)
47
Seven Planetary Elements of Alchemists
48
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Other Alchemical Elements
As
Sb
Bi
Pt
S
P
50
Law of Triads (Dobereiner, 1829)
The molar mass and density of the middle one ?
average of the other two.
51
Law of Octaves (Newlands, 1865)
Elements of similar physical and chemical
properties recurred at intervals of eight
Li Be B C N O F
Na Mg Al Si P S Cl
52
First Periodic Table (Mendeleev, 1869)
Periodicity Chemical properties of elements are
periodic functions of their atomic
masses. Elements arranged in terms of their
properties (not exactly follow the order of
atomic mass) Elements with similar properties
are put together in vertical groups Gaps were
left in the table for missing elements
53
First Periodic Table (Mendeleev, 1869)

• missing elements predicted by Mendeleev
• Ekaboron (atomic mass 44)
• Scandium (44.96)
• Ekaaluminium (68)
• Gallium (69.3)

54
First Periodic Table (Mendeleev, 1869)
missing elements predicted by Mendeleev

• 3. Ekamanganese (100)
• Technetium (98)
• Ekasilicon (72)
• Germanium (72.59)

55
7 groups or 8 groups ?
56
Discovery of the Noble Gases
Lord Rayleigh
William Ramsay
Nobel Laureate in Physics, 1904
Nobel Laureate in Chemistry, 1904
57
1894
Density ( g / dm?3)
1.2572
decompose
1.2508
error ? 0.5
???
58

• Argon is present in air
• Confirmed by spectroscopy
• RAM Ar(39.95) gt K(39.10)
• Unlike group 2 elements, Ar shows no reactivity.
• Placed before K and after Cl
• A new group in the Periodic Table
• Group 0

59
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar
Rn discovered in 1900 by F.E. Dorn
Po, Ra discovered in 1898 by Pierre Marie Curie
60
Congratulations ! Nobel Laureate in
Chemistry, 2010
61
Modern Periodic Table
Elements arranged in order of increasing atomic
number
91 elements discovered up to 1940 Most are
naturally occurring except Po(84), At(85),
Rn(86), Fr(87), Ra(88), Ac(89), Pa(91) from
radioactive decay Pm(61) discovered in 1945 as a
product in nuclear fission - not found in nature
62
Transuranium Elements
92 U
Discovered by McMillan and Seaborg
63
Transuranium Elements
92 U 93 Np 94 Pu
Discovered by McMillan and Seaborg
Nobel Laureates in Chemistry, 1951 From
University of California, Berkeley, United States
of America
64
92 U 93 94 95 96 94 96 98 99
Uranium, discovered in 1789, was considered the
heaviest elements
65
92 U
Named after Uranus (???) Discovered in 1781 Was
Considered the Farthest Planet from The Earth in
the Solar System
66
Transuranium Elements
92 U 93 Np 94 Pu
Neptunium Discovered in 1940 by McMillan
Neptune(???) The Next Planet out from Uranus
67
Transuranium Elements
92 U 93 Np 94 Pu
Plutonium Discovered in 1941 by
McMillan Seaborg
Pluto(???) Was considered the next Planet out
from Neptune
68
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Americium (1944)
Nobel Laureates in Chemistry, 1951 University of
California, Berkeley, United States of America
69
It was named americium because it is just below
europium in the Periodic Table.
70
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Curium
Marie Curie
Nobel Laureate in Physics, 1903 Nobel Laureate in
Chemistry, 1911
71
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Berkelium (1949)
Nobel Laureates in Chemistry, 1951 University of
California, Berkeley, United States of America
72
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Californium (1950)
Nobel Laureates in Chemistry, 1951 University of
California, Berkeley, United States of America
73
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
McMillan and Seaborg
Nobel Laureates in Chemistry, 1951
74
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Einsteinium (1952 by Albert Ghiorso)
Albert Einstein
Nobel Laureate in Physics, 1921
75
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Fermium (1952 by Albert Ghiorso)
Enrico Fermi
Nobel Laureate in Physics, 1938 Developer of the
first nuclear reactor, 1942
76
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
General Consultant of the Manhattan
Project Hiroshima little boy Nagasaki fat man
77
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Mendelevium (1955 by Albert Ghiorso)
Mendeléev
Discovery of Periodicity 1869
78
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Nobelium (1958 by Albert Ghiorso)
Alfred Nobel
Inventor of Dynamite, 1867 The Man Behind the
Nobel Prize
79
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Lawrencium (1961 by Albert Ghiorso)
_at_ Lawrence Radiation Laboratory
University of California, Berkeley,
80
Transuranium Elements
95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr
Ernest Orlando Lawrence
Developer of cyclotron
Nobel Laureate in Physics, 1939
University of California, Berkeley,
81
1964 1996 AD Teams from Russia(USSR), USA
Germany Synthesis of Rf(104), Db(105), Sg(105),
Bh(106), Hs(108), Mt(109), Ds(110), Rg(111)
Uub(112).
Rg(111) Roentgenium
82
1999 2003 AD Russia(USSR) USA Synthesis of
Uut(113), Uuq(114), Uup(115), Uuh(116)
83
Naming of Elements IUPAC System
111 unununium (Uuu) Roentgenium (Rg)
84
111 unununium (Uuu)
d-block
112 ununbium (Uub)
113 ununtrium (Uut)
114 ununquadium (Uuq)
p-block
115 ununpentium (Uup)
116 ununhexium (Uuh)
6B
85
The Periodic Table
5.3 The Periodic Table (SB p.112)
86
5.3 The Periodic Table (SB p.112)
s-block p-block elements are called
representative elements
s-block
p-block
d-block
f-block
87
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88

• ?????
• ?????
• ???

89

• ?(La)??(Ce)??(Pr)??(Nd)??(Pm)?
• ?(Sm)??(Eu)??(Gd)??(Tb)??(Dy)?
• ?(Ho)??(Er)??(Tm)??(Yb)??(Lu)?
• ?(Sc)??(Y)

17 Rare Earth Metals(????)
90
Q.23(a)
They are named after the outermost orbitals to be
filled
91
Q.23(b)
No
d- block f-block
Period no. n n
No. of the last subshell to be filled n 1 n ? 4 n 2 n ? 6
92
d-block starts in Period 4 (n ? 4) Transition
metals f-block starts in Period 6 (n ?
6) Lanthanides Period 6
(rare earth metals) Actinides Period 7
93
Q.23(c)
True only for IB to VII B
94
Q.23(c)
IIIB Sc Ar 3d1, 4s2 IVB Ti Ar 3d2,
4s2 VB V Ar 3d3, 4s2 VIB Cr Ar 3d5,
4s1 VIIB Mn Ar 3d5, 4s2
95
Q.23(c)
IB Cu Ar 3d10, 4s1 IIB Zn Ar 3d10, 4s2
Electrons in fully-filled 3d subshells cannot be
removed easily. ? They are not treated as
outermost shell electrons
96
Q.23(c)
Not true for VIIIB elements
VIIIB Fe Ar 3d6, 4s2 Co Ar 3d7,
4s2 Ni Ar 3d8, 4s2
97
5.3 The Periodic Table (SB p.112)
Let's Think 1
98
The Song of Elements by Tom Lehrer
The Song of Elements on YouTube
Visual Elements Periodic Table
99

• Periodicity as illustrated by
• Variation in atomic radius with atomic number
• Variation in ionization enthalpy with atomic
  number

100
Atomic Size of Elements
101
Atomic radius is defined as half the distance
between two nuclei of the atoms joined by a
single covalent bond or a metallic
bond
102
Atomic radii of noble gases were obtained by
calculation
103
Atomic radii ? across both Periods 2 and 3
104
5.6 Atomic size of elements (p. 122)
Q Explain why the atomic radius decreases across
a period.

• Moving across a period, there is an increase in
  the nuclear attraction due to the addition of
  proton in the nucleus.(? in nuclear charge)
• The added electron is placed in the same quantum
  shell. It is only poorly repelled/shielded/screen
  ed by other electrons in that shell.
• The nuclear attraction outweighs the increase in
  the shielding effect between the electrons. This
  leads to an increase in the effective nuclear
  charge.

105
Effective nuclear charge, Zeff, is the
nuclear charge experienced by an electron in an
atom. In the present discussion, only the
outermost electrons are considered.
106
3
Li
The outer 2s electron sees the nucleus through a
screen of two inner 1s electrons.
107
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108
The outer 2s electron is repelled/shielded/screene
d by the inner 1s electrons from the nucleus
109
The nuclear charge experienced by the 2s electron
is ? 1
110
4
Be
The inner 1s electrons shield the outer electrons
almost completely
111
2
Be
The two electrons in the same shell (2s) shield
each other less poorly.
Zeff ? 1.5
112
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113
Atomic radii ? down a group
114
5.6 Atomic size of elements (p. 122)
Q Explain why the atomic radius increases down a
group.

• Moving down a group, an atom would have more
  electron shells occupied. The outermost shell
  becomes further away from the nucleus.
• Moving down a group, although there is an
  increase in the nuclear charge, it is offset very
  effectively by the screening effect of the inner
  shell electrons.

115
Sharp ? in atomic radius when a new Period begins
116
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is sharp ? in atomic radius
when a new Period begins

• The element at the end of a period has the
  smallest atomic radius among the elements in the
  same period because its outermost electrons are
  experiencing the strongest nuclear attraction.

117
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is sharp ? in atomic radius
when a new Period begins

• The element at the beginning of the next period
  has one extra electron in an outer shell which is
  far away from the nucleus. Although there is also
  an increase in the nuclear charge, it is very
  effectively screened by the inner shell electrons.

118
Ionization Enthalpies of Elements
119
Across a Period, there is a general ? in I.E.
leading to a maximum with a noble gas.
120
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121
First I.E. ? down a group
122
The outermost electrons are further away from the
nucleus and are more effectively shielded from it
by the inner electrons
123
5.4 Ionization enthalpies of elements (SB p.116)
The first ionization enthalpies generally
decrease down a group and increase across a period
124
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125
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is sharp ? in IE when a
new Period begins

• The element at the end of a period has a stable
  duplet or octet structure. Much energy is
  required to remove an electron from it as this
  will disturb the stable structure.

126
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is sharp ? in I.E. when a
new Period begins

• The element at the beginning of the next period
  has one extra electron in an outer quantum shell
  which is far away from the nucleus.
• Although there is also an increase in the nuclear
  charge, it is very effectively shielded by the
  inner shell electrons.
• Thus the outermost electron experiences a much
  less nuclear attraction.

127
Irregularities -
128
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is a trough at Boron(B) in
Period 2.

• Be 1s2, 2s2B 1s2, 2s2, 2p1

129
More diffused
130
In multi-electron systems, penetrating power
- s gt p gt d gt f
131
3d electrons are more diffused (less
penetrating) 3d electrons are more shielded by 1s
electrons
132
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is a trough at Boron(B) in
Period 2.

• It is easier to remove the less penetrating 2p
  electron from B than to remove a more penetrating
  2s electron from a stable fully-filled 2s
  subshell in Be.

133
Irregularities -
134
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is a trough at Oxygen(O) in
Period 2.

• e.c. of N 1s2, 2s2, 2px1, 2py1, 2pz1e.c. of
  O 1s2, 2s2, 2px2, 2py1, 2pz1
• The three 3p electrons in N occupy three
  different orbitals, thus minimizing the repulsion
  between the electrons(shielding effect). It is
  more difficult to remove an electron from the
  half-filled 2p subshell of N.

135
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is a trough at Oxygen(O) in
Period 2.

• e.c. of N 1s2, 2s2, 2px1, 2py1, 2pz1
• e.c. of O 1s2, 2s2, 2px2, 2py1, 2pz1
• Alternately, the removal of a 2p electron
• from O results in a stable half-filled 2p
• subshell.

136
Variation of Successive Ionization Enthalpies
with Atomic Numbers
137
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 120)
Successive I.Es. Show similar variation patterns
with atomic number. 3rd I.E. gt 2nd I.E. gt 1st I.E.
138
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 120)
Plots of successive I.E. are shifted by one unit
in atomic number to the right respectively. e.g. B
e Li (2, 1) B Be (2, 2) C B
(2, 3)
Each represents a pair of isoelectronic species
139
Invert relationship between atomic radius and
first I.E.
Why is the atomic radius of helium greater than
that of hydrogen, despite of the fact that the
first I.E. of helium is higher than that of
hydrogen ?
140
Q.24

• (a) A would have the largest atomic number.
• It is because A has the lowest first
  ionization enthalpy.
• Group I
• It is because 1st I.E. ltlt 2nd I.E.

141
Q.25

• B is most likely to form B3
• It is because 3rd I.E. ltlt 4th I.E.
• A and D are in Group I
• It is because 1st I.E. ltlt 2nd I.E.

142
Q.26

• D is a noble gas.
• It is because D has a higher I.E. than those
  of A, B and C and has a much higher I.E. than E.
• A B C D E F
• N O F Ne Na Mg
• P S Cl Ar K
  Ca

143
The END
144
5.1 Relative energies of orbitals (SB p.108)
Back
Check Point 5-1

• Write the electronic configurations and draw
  electrons-in boxes diagrams for
• (a) nitrogen and
• (b) sodium.

Answer
145
5.2 Electronic configurations of elements (SB
p.110)
Back
Check Point 5-2

• Give the electronic configuration by notations
  and electrons-in-boxes diagrams in the
  abbreviated form for the following elements.
• silicon and
• copper.

Answer
146
5.3 The Periodic Table (SB p.113)
Back
Let's Think 1
If you look at the Periodic Table in Fig. 5-5
closely, you will find that hydrogen is separated
from the rest of the elements. Even though it has
only one electron in its outermost shell, it
cannot be called an alkali metal, why?
Answer
Hydrogen has one electron shell only, with n 1.
This shell can hold a maximum of two electrons.
Hydrogen is the only element with core electrons.
This gives it some unusual properties. Hydrogen
can lose one electron to form H, or gain an
electron to become H-. Therefore, it does not
belong to the alkali metals and halogens.
Hydrogen is usually assigned in the space above
the rest of the elements in the Periodic Table
the element without a family.
147
5.3 The Periodic Table (SB p.114)
Check Point 5-3
Outline the modern Periodic Table and label the
table with the following terms representative
elements, d-transition elements, f-transition
elements, lanthanide series, actinide series,
alkali metals, alkaline earth metals, halogens
and noble gases.
Answer
148
5.3 The Periodic Table (SB p.114)
Back
Check Point 5-3
149
5.4 Ionization enthalpies of elements (SB p.118)
Check Point 5-4

• Give four main factors that affect the magnitude
  of ionization enthalpy of an atom.

Answer

• The four main factors that affect the magnitude
  of the ionization enthalpy of an atom are
• (1) the electronic configuration of the atom
• (2) the nuclear charge
• (3) the screening effect and
• (4) the atomic radius.

150
5.4 Ionization enthalpies of elements (SB p.118)
Check Point 5-4

• Explain why Group 0 elements have extra high
  first ionization enthalpies and their decreasing
  trend down the group.

Answer

• The first ionization enthalpies of Group 0
  elements are extra high. It is because Group 0
  elements have very stable electronic
  configurations since their orbitals are
  completely filled. That means, a large amount of
  energy is required to remove an electron from a
  completely filled electron shell of ns2np6
  configuration.
• Going down the group, the first ionization
  enthalpies of Group 0 elements decreases. It is
  because there is an increase in atomic radius
  down the group, the outermost shell electrons
  experience less attraction from the nucleus.
  Further, as there is an increase in the number of
  inner electron shells, the outermost shell
  electrons of the atoms are better shielded from
  the attraction of the nucleus (greater screening
  effect). Consequently, though the nuclear charge
  increases down the group, the outermost shell
  electrons would experience less attraction from
  the positively charged nucleus. That is why the
  first ionization enthalpies decrease down the
  group.

151
5.4 Ionization enthalpies of elements (SB p.118)
Back
Check Point 5-4

• Predict the trend of the first ionization
  enthalpies of the transition elements.

Answer
(c) The first ionization enthalpies of the
transition elements do not show much variation.
The reason is that the first electron of these
atoms to be removed is in the 4s orbital. As the
energy levels of the 4s orbitals of these atoms
are more or less the same, the amount of energy
required to remove these electrons are similar.
152
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 121)
Example 5-5

• For the element 126C,
• (i) write its electronic configuration by
  notation.
• (ii) write its electronic configuration by
  electrons-in- boxes diagram.

Answer
153
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 121)
Example 5-5

• The table below gives the successive ionization
  enthalpies of carbon.
• (i) Plot a graph of log ionization enthalpy
  against number of electrons removed.
• (ii) Explain the graph obtained.

1st 2nd 3rd 4th 5th 6th
I.E. (kJ mol-1) 1090 2350 4610 6220 37800 47000
Answer
154
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 121)
Example 5-5
155
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 121)
Back
Example 5-5
(ii) The ionization enthalpy increases with
increasing number of electrons removed. It is
because the effective nuclear charge increases
after an electron is removed, and more energy is
required to remove an electron from a positively
charged ion. Besides, there is a sudden rise
from the fourth to the fifth ionization
enthalpy. This is because the fifth ionization
enthalpy involves the removal of an electron
from a completely filled 1s orbital which is
very stable.
156
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 122)
Check Point 5-5

1. Give the electrons-in-boxes diagram of 26Fe.
2. Fe2 and Fe3 have 2 and 3 electrons less than Fe
   respectively. If the electrons are removed from
   the 4s orbital and then 3d orbitals, give the
   electronic configurations of Fe2 and Fe3.

Answer
157
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 122)
Check Point 5-5

• (c) Which ion is more stable, Fe2 or Fe3?
  Explain briefly.

(c) Fe3 ion is more stable because the 3d
orbital is exactly half-filled which gives the
electronic configuration extra stability.
Answer
158
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 122)
Check Point 5-5

• Given the successive ionization enthalpies of Fe
• (i) plot a graph of successive ionization
  enthalpies in logarithm scale against the
  number of electrons removed
• (ii) state the difference of the plot from that
  of carbon as shown in P. 121.

1st 2nd 3rd 4th 5th 6th
I.E. (kJ mol-1) 762 1560 2960 5400 7620 10100
Answer
159
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 122)
Check Point 5-5

• (i)

Number of electrons removed 1 2 3 4 5 6
log (I.E.) 2.88 3.19 3.47 3.73 3.88 4.00
160
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 122)
Check Point 5-5
(ii) The ionization enthalpy increases with
increasing number of electrons removed. This is
because it requires more energy to remove an
electron from a higher positively charged ion. In
other words, higher successive ionization
enthalpies will have higher magnitudes. However
, the sudden increase from the fourth to the
fifth ionization enthalpies occurs in carbon but
not in iron. This indicates that when electrons
are removed from the 4s and 4d orbitals, there
is no disruption of a completely filled electron
shell. Hence, there are no irregularities for
the first six successive ionization enthalpies
of iron.
Back
161
5.6 Atomic size of elements (p. 123)
Check Point 5-6

• Explain the following
• (a) The atomic radius decreases across the
  period from Li to Ne.

Answer
(a) When moving across the period from Li to Ne,
the atomic sizes progressively decrease with
increasing atomic numbers. This is because an
increase in atomic number by one means one more
electron and one more proton in atoms. The
additional electron would cause an increase in
repulsion between the electrons in the outermost
shell. However, since each additional electron
goes to the same quantum shell and is at
approximately the same distance from the nucleus,
the repulsion between electrons is relatively
ineffective to cause an increase in the atomic
radius. On the other hand, as there is an
additional proton added to the nucleus, the
electrons will experience a greater attractive
force from the nucleus (increased effective
nuclear charge). Hence, the atomic radii of atoms
decrease across the period from Li to Ne.
162
5.6 Atomic size of elements (p. 123)
Back
Check Point 5-6

• Explain the following
• (b) The atomic radius increases down Group I
  metals.

Answer
(b) Moving down Group I metals, the atoms have
more electron shells occupied. The outermost
electron shells become further away from the
nucleus. Besides, the inner shell electrons will
shield the outer shell electrons more effectively
from the nuclear charge. This results in a
decrease in the attractive force between the
nucleus and the outer shell electrons. Therefore,
the atomic radii of Group I atoms increase down
the group.