5
5.1
5.2
Relative Energies of Orbitals
Electronic Configurations of Elements
5.3
5.4
5.5
The Periodic Table
Ionization Enthalpies of Elements
Variation of Successive Ionization Ethalpies
with Atomic Numbers
Atomic Size of Elements
5.4
1
Electronic Configurations
and the Periodic Table
5.1
Relative Energies
of Orbitals
2
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.
3
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
2s
2p
1s
4
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 < p < d < f
5
5.1 Relative energies of orbitals (SB p.106)
Relative energies of orbitals
6
5.1 Relative energies of orbitals (SB p.106)
Relative energies of orbitals
7
5.1 Relative energies of orbitals (SB p.106)
Relative energies of orbitals
8
5.1 Relative energies of orbitals (SB p.106)
Relative energies of orbitals
Electrons enter 4s subshell
before filling up 3d subshell.
9
Both 4s and 3d electrons are shielded from the
nuclear attraction by the inner core (2,8,8)
10
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.
11
Three rules to build up electronic
configurations
1. Aufbau (building up) Principle
2. Hund’s Rule
3. Pauli’s Exclusion Principle
12
1. Aufbau (building up) Principle
Electrons enter the orbitals in
order of ascending energy.
13
Numbers
Letters read
read downwards
across
s
p
d
f
g
h
i
1
14
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s,
s
p
d
f
g
h
i
1
15
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s, 2s,
s
p
d
f
g
h
i
1
16
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s, 2s, 2p, 3s,
s
p
d
f
g
h
i
1
17
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s, 2s, 2p, 3s, 3p, 4s,
s
p
d
f
g
h
i
1
18
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s
s
p
d
f
g
h
i
1
19
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s
s
p
d
f
g
h
i
1
20
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s
s
p
d
f
g
h
i
1
21
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
7
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
22
2
2
3
3
3
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
7
7
7
7
7
7
8
7
5.1 Relative energies of orbitals (SB p.106)
Building up of electronic configurations
23
2. Hund’s rule : -
Orbitals of the same energy must be
occupied singly and with the same spin
before pairing up of electrons occurs.
Carbon
1s
2s
2p
Electrons-in boxes diagram
24
3.
Pauli’s exclusion principle : -
Electrons occupying the same
orbital must have opposite spins.
W. Pauli , Nobel prize
laureate in Physics, 1945
Check Point 5-1
25
5.2
Electronic
Configurations of
Elements
26
5.2 Electronic configurations of elements (SB p.108)
Ways to Express Electronic Configurations
1. The s, p, d, f notation
Na 1s2, 2s2, 2p6, 3s1
1s2, 2s2, 2px2, 2py2, 2pz2, 3s1
27
Q.17(a)
K
1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Q.17(b)
Fe
1s2, 2s2, 2p6, 3s2, 3p6, 3d6, 4s2
orbitals of the same quantum shell
are placed together
28
2.
Using a noble gas ‘core’
Na [Ne] 3s1
Ca [Ar] 4s2
29
Q.18(a)
Si
[Ne] 3s2, 3p2
outermost shell
Q.18(b)
V
[Ar] 3d3, 4s2
outermost shell
30
3.
Electrons – in – Boxes representation
N
1s
2s
2p
(1) All boxes should be labelled
(2) Boxes of the same energies are put
together.
2px
31
2py
2pz

Q.19
Hund’s rule is violated
Hund’s rule is violated
Pauli’s exclusion principle is violated
32
Q.20(a)
Phosphorus
1s
33
2s
2p
3s
3p
Q.20(b)
Chromium
1s
2s
2p
3s
3p
3d
4s
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.
34
Q.20(b)
[Ar] 3d4, 4s2  [Ar] 3d5, 4s1 + energy
Chromium
1s
2s
2p
3s
3p
3d
4s
1s
2s
2p
3s
3p
3d
4s
+
energy
35
Q.20(c)
Copper
1s
2s
2p
3s
3p
3d
4s
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.
36
Q.20(c)
[Ar] 3d9, 4s2  [Ar] 3d10, 4s1 + energy
Copper
1s
2s
2p
3s
3p
3d
4s
1s
2s
2p
3s
3p
3d
4s
+
energy
37
Silicon
3s
3p
[Ne]
Empty orbital(s) in a partially filled
subshell should be shown
38
Silicon
3s
3p
3s
3p
[Ne]
[Ne]
+ energy

Energy difference : 3p – 3s > 3d – 4s
39
21(a)
21(b)
Ar
(i) S2
1s2, 2s2, 2p6, 3s2, 3p6
(ii) Cl
1s2, 2s2, 2p6, 3s2, 3p6
(iii) K+
1s2, 2s2, 2p6, 3s2, 3p6
(iv) Ca2+ 1s2, 2s2, 2p6, 3s2, 3p6
40
S2 , Cl , Ar , K+ , Ca2+
Same electronic configurations
Isoelectronic
Q.22
41
5.2 Electronic configurations of elements (SB p.110)
Represented by ‘electrons-in-boxes’ diagrams
42
5.2 Electronic configurations of elements (SB p.110)
43
Check Point 5-2
Building up of electronic configurations
http://www.chemcollective.org/applets/
pertable.php
44
A brief history of the Periodic Table
Ancient Greece,
Aristotle : - Four elements
Fire,
45
Water,
Air,
Earth,
A brief history of the Periodic Table
Ancient Greece,
Aristotle : - Four elements
Air, Fire, Earth, Water
Buddha : 地、水、火、風、空
Quintessence (The fifth element)
46
Seven Planetary Elements of Alchemists
47
Moon  Silver
Mars  Iron
Sun  Gold
Venus  Copper
Jupiter  Tin
48
Mercury  Mercury
Saturn  Lead
Other Alchemical Elements
As
Sb
Pt
P
49
Bi
S
Law of Triads (Dobereiner, 1829)
Element Molar mass Density
(g/mol)
(g/cm³)
50
chlorine
35.453
0.0032
bromine
79.904
3.1028
iodine
126.90447
4.933
calcium
40.078
1.55
strontium 87.62
2.54
barium
3.594
137.327
The molar mass and
density of the middle
one  average of the
other two.
Law of Octaves (Newlands, 1865)
Elements of similar physical and chemical
properties recurred at intervals of eight
Group
1A
51
Group Group
2A
3A
Group Group Group
4A
5A
6A
Group
7A
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
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’
52
First Periodic Table (Mendeleev, 1869)
‘missing elements’ predicted by Mendeleev
1. Ekaboron (atomic mass = 44)
Scandium (44.96)
2. Ekaaluminium (68)
Gallium (69.3)
53
First Periodic Table (Mendeleev, 1869)
‘missing elements’ predicted by Mendeleev
3. Ekamanganese (100)
Technetium (98)
4. Ekasilicon (72)
Germanium (72.59)
54
7 groups or 8 groups ?
55
Discovery of the Noble Gases
Lord Rayleigh
Nobel Laureate in
Physics, 1904
56
William Ramsay
Nobel Laureate in
Chemistry, 1904
1894
Air
Density ( g / dm3)
- (O2, CO2, H2O)
NH3
decompose
N2
1.2572
N2
1.2508
% error  0.5%
57
???
Argon is present in air
Confirmed by spectroscopy
RAM : Ar(39.95) > 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
58
Group
1A
Group Group Group Group Group Group Group
2A
3A
4A
5A
6A
7A
0
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Helium discovered in 1895
Ne, Kr, Xe discovered in 1898
All by
Ramsay
Rn discovered in 1900 by F.E. Dorn
Po, Ra discovered in 1898 by Pierre & Marie Curie
59
Congratulations !
Nobel Laureate in Chemistry,
2010
60
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
61
Transuranium Elements
92
U
93
Np
94
95
96
94
96
98
Pu
Discovered by McMillan and Seaborg
62
99
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
63
92
93
94
95
96
94
96
98
99
U
Uranium, discovered in 1789,
was considered the heaviest elements
64
92
U
Named after Uranus (天王星)
Discovered in 1781
Was Considered the Farthest Planet
from The Earth in the Solar System
65
Transuranium Elements
92
U
93
Np
94
Pu
Neptunium : Discovered in 1940 by
McMillan
Neptune(海王星) : The Next Planet
out from Uranus
66
Transuranium Elements
92
U
93
Np
94
Pu
Plutonium : Discovered in 1941 by
McMillan & Seaborg
Pluto(冥王星) : Was considered the
next ‘Planet’ out from Neptune
67
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
Americium (1944)
Nobel Laureates in Chemistry, 1951
University of California, Berkeley,
United States of America
68
103
Lr
It was named americium because it is
just below europium in the Periodic
Table.
69
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
Curium
100
101
102
Fm Md No
Marie Curie
Nobel Laureate in Physics, 1903
Nobel Laureate in Chemistry, 1911
70
103
Lr
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
Berkelium (1949)
Nobel Laureates in Chemistry, 1951
University of California, Berkeley,
United States of America
71
103
Lr
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
Californium (1950)
Nobel Laureates in Chemistry, 1951
University of California, Berkeley,
United States of America
72
103
Lr
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
McMillan and Seaborg
Nobel Laureates in Chemistry, 1951
73
103
Lr
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
103
Lr
Einsteinium (1952 by Albert Ghiorso)
Albert Einstein
Nobel Laureate in Physics, 1921
74
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
103
Lr
Fermium (1952 by Albert Ghiorso)
Enrico Fermi
Nobel Laureate in Physics, 1938
Developer of the first nuclear
reactor, 1942
75
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
Nagasaki – fat man
76
102
Fm Md No
General Consultant of the
Manhattan Project
Hiroshima – little boy
101
103
Lr
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
103
Lr
Mendelevium (1955 by Albert Ghiorso)
Mendeléev
Discovery of Periodicity 1869
77
Transuranium Elements
95
96
97
98
Am Cm Bk
Cf
99
Es
100
101
102
Fm Md No
103
Lr
Nobelium (1958 by Albert Ghiorso)
Alfred Nobel
Inventor of Dynamite, 1867
The Man Behind the Nobel Prize
78
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
102
Fm Md No
103
Lr
Lawrencium (1961 by Albert Ghiorso)
@ Lawrence Radiation Laboratory
University of California, Berkeley,
79
Transuranium Elements
95
96
97
Am Cm Bk
98
Cf
99
Es
100
101
Fm Md No
Ernest Orlando Lawrence
Developer of cyclotron
Nobel Laureate in Physics, 1939
University of California,
Berkeley,
80
102
103
Lr
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
81
1999 – 2003 AD
Russia(USSR) & USA
Synthesis of
Uut(113), Uuq(114), Uup(115), Uuh(116)
82
Naming of Elements – IUPAC System
0 = nil
1 = un
2 = bi
3 = tri
4 = quad
ium
5 = pent
6 = hex
7 = sept
8 = oct
9 = enn
ium
111 = unununium (Uuu) = Roentgenium (Rg)
83
111 = unununium (Uuu)
112 = ununbium (Uub)
d-block
113 = ununtrium (Uut)
114 = ununquadium (Uuq)
115 = ununpentium (Uup)
p-block
116 = ununhexium (Uuh)
6B
84
5.3 The Periodic Table (SB p.112)
The Periodic Table
85
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
86
Rare Earth Metals(稀土金屬)
f-block elements are called
inner-transition elements
87
中東有石油
中國有稀土
鄧小平
88
鑭(La)、鈰(Ce)、鐠(Pr)、釹(Nd)、鉕(Pm)、
釤(Sm)、銪(Eu)、釓(Gd)、鋱(Tb)、鏑(Dy)、
鈥(Ho)、鉺(Er)、銩(Tm)、鐿(Yb)、鑥(Lu)、
鈧(Sc)、釔(Y)
17 Rare Earth Metals(稀土金屬)
89
Q.23(a)
They are named after the outermost
orbitals to be filled
90
Q.23(b)
91
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
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
92
Q.23(c)
True only for IB to VII B
93
Q.23(c)
94
IIIB
Sc
[Ar] 3d1, 4s2
3
IVB
Ti
[Ar] 3d2, 4s2
4
VB
V
[Ar] 3d3, 4s2
5
VIB
Cr
[Ar] 3d5, 4s1
6
VIIB
Mn
[Ar] 3d5, 4s2
7
Q.23(c)
IB
Cu
[Ar] 3d10, 4s1
IIB
Zn
[Ar] 3d10, 4s2
Electrons in fully-filled 3d subshells
cannot be removed easily.

95
They are not treated as outermost
shell electrons
Q.23(c)
Not true for VIIIB elements
VIIIB
96
Fe
[Ar] 3d6, 4s2
Co
[Ar] 3d7, 4s2
Ni
[Ar] 3d8, 4s2
5.3 The Periodic Table (SB p.112)
Check Point 5-3
97
The Song of Elements by Tom Lehrer
The Song of Elements – on YouTube
Visual Elements Periodic Table
98
Periodicity as illustrated by
(i) Variation in atomic radius with
atomic number
(ii) Variation in ionization enthalpy
with atomic number
99
5.6
Atomic Size of
Elements
100
Atomic radius is defined as
half the distance between two
nuclei of the atoms joined by a
single covalent bond or
a metallic bond
101
Atomic radii of noble gases were obtained by
calculation
102
Atomic radii  across both Periods 2 and 3
103
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/screened 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.
104
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.
105
+3
Li
The outer 2s electron sees the nucleus
through a screen of two inner 1s
electrons.
106
Two electrons in
the inner shell
nucleus
107
2s electron of
Li outside the
inner shell
+3
Li
The outer 2s electron is
repelled/shielded/screened by the
inner 1s electrons from the nucleus
108
+3
Li
 +1
Li
The nuclear charge experienced by the
2s electron is  +1
109
+4
+2
Be
Be
The inner 1s electrons shield the outer
electrons almost completely
110
+2
Be
+1.5
Be
The two electrons in the same shell (2s)
shield each other less poorly.
Zeff  1.5
111
112
Atomic radii  down a group
113
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.
114
Sharp  in atomic radius when a new Period begins
115
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.
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 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.
Check Point 5-6
117
5.4
Ionization Enthalpies
of Elements
118
Across a Period, there is a general  in I.E.
leading to a maximum with a noble gas.
119
 Effective nuclear charge  from left to
right across the Period
120
First I.E.  down a group
121
The outermost electrons are
further away from the nucleus
and are more effectively shielded
from it by the inner electrons
122
5.4 Ionization enthalpies of elements (SB p.116)
The first ionization enthalpies
generally decrease down a group
and increase across a period
123
124
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.
He  1s2 (duplet)
Ne  2s2, 2p6 (octet)
Ar  3s2, 3p6 (octet)
125
Fully-filled
shells
Fully-filled
subshell
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.
126
Irregularities : -
127
5.4 Ionization enthalpies of elements (SB p.117)
Q: Explain why there is a trough at Boron(B)
in Period 2.
• Be : 1s2, 2s2
B : 1s2, 2s2, 2p1
128
More
penetrating
1s
More
diffused
129
In multi-electron
systems,
penetrating power : -
s>p>d>f
130
 2 4r 2
3d electrons are more diffused
(less penetrating)
3d electrons are more shielded by
1s electrons
131
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.
132
Irregularities : -
133
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
• 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.
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, 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.
135
5.5
Variation of Successive
Ionization Enthalpies
with Atomic Numbers
136
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. > 2nd I.E. > 1st I.E.
137
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+ = Be
C+ = B
138
Each represents
a pair of
(2, 2)
isoelectronic
(2, 3)
species
Be+ = Li (2, 1)
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 ?
139
Q.24
(a)
A would have the largest atomic
number.
It is because A has the lowest first
ionization enthalpy.
(b) Group I
It is because 1st I.E. << 2nd I.E.
140
Q.25
(a)
B is most likely to form B3+
It is because 3rd I.E. << 4th I.E.
(b) A and D are in Group I
It is because 1st I.E. << 2nd I.E.
141
Q.26
(a)
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.
(b) A
B
C
D
E
F
N
O
F
Ne
Na
Mg
P
S
Cl
Ar
K
Ca
142
The END
143
5.1 Relative energies of orbitals (SB p.108)
Back
Write the electronic configurations and draw “electrons-in –
boxes” diagrams for
(a) nitrogen; and
(b) sodium.
(a) Nitrogen: 1s22s22p3
(b) Sodium: 1s22s22p63s1
144
Answer
5.2 Electronic configurations of elements (SB p.110)
Back
Give the electronic configuration by notations and
“electrons-in-boxes” diagrams in the abbreviated form for
the following elements.
(a) silicon; and
(b) copper.
(a) Silicon: [Ne]3s23p3
(b) Copper: [Ar]3d104s1
145
Answer
5.3 The Periodic Table (SB p.113)
Back
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.
146
5.3 The Periodic Table (SB p.114)
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
147
5.3 The Periodic Table (SB p.114)
Back
148
5.4 Ionization enthalpies of elements (SB p.118)
(a) Give four main factors that affect the magnitude of
ionization enthalpy of an atom.
Answer
(a) 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.
149
5.4 Ionization enthalpies of elements (SB p.118)
(b) Explain why Group 0 elements have extra high first
ionization enthalpies and their decreasing trend down the
group.
(b) The first ionization enthalpies of Group 0 elements are extra high. It is
Answer
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
150positively charged nucleus. That is why the first ionization enthalpies
decrease down the group.
5.4 Ionization enthalpies of elements (SB p.118)
Back
(c) 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.
151
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
For the element 126C,
(a) (i) write its electronic configuration by notation.
(ii) write its electronic configuration by “electrons-inboxes” diagram.
(a) (i) 1s22s22p2
(ii)
152
Answer
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
(b) The table below gives the successive ionization
enthalpies of carbon.
I.E. (kJ
mol-1)
1st
1090
2nd
2350
3rd
4610
4th
6220
5th
37800
6th
47000
(i) Plot a graph of log [ionization enthalpy] against
number of electrons removed.
(ii) Explain the graph obtained.
153
Answer
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
(b) (i)
154
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
Back
(ii)
155
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.
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(a) Give the “electrons-in-boxes” diagram of 26Fe.
(a) Fe :
(b) 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+.
(b) Fe2+ :
Fe3+ :
156
Answer
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(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
157
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(d) Given the successive ionization enthalpies of Fe:
I.E. (kJ
mol-1)
1st
762
2nd
1560
3rd
2960
4th
5400
5th
7620
6th
10100
(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.
158
Answer
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(d) (i)
159
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
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(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
160
5.6 Atomic size of elements (p. 123)
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
161 atoms decrease across the period from Li to Ne.
5.6 Atomic size of elements (p. 123)
Back
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.
162