1 electronic configurations and the periodic table 5.1relative energies of orbitals 5.2electronic...
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1
Electronic Configurations Electronic Configurations
and the Periodic Tableand the Periodic Table
5.15.1 Relative Energies of Orbitals Relative Energies of Orbitals
5.25.2 Electronic Configurations of Elements Electronic Configurations of Elements
5.35.3 The Periodic TableThe Periodic Table
5.45.4 Ionization Enthalpies of ElementsIonization Enthalpies of Elements5.55.5 Variation of Successive Ionization EthalpiesVariation of Successive Ionization Ethalpies with Atomic Numbers with Atomic Numbers
5.45.4 Atomic Size of ElementsAtomic Size of Elements
55
2
5.5.11Relative EnergieRelative Energie
s of Orbitalss 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
In the Lyman series,
only one spectral line is observed for the transition from n = 2 to n = 1.
Evidence
2s 2p
1s
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 < p < d < f
6
Relative energies of orbitalsRelative energies of orbitals5.1 Relative energies of orbitals (SB p.106)
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Relative energies of orbitalsRelative energies of orbitals5.1 Relative energies of orbitals (SB p.106)
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Relative energies of orbitalsRelative energies of orbitals5.1 Relative energies of orbitals (SB p.106)
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Relative energies of orbitalsRelative energies of orbitals5.1 Relative energies of orbitals (SB p.106)
Electrons enter 4s subshell before filling up 3d subshell.
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. Hund’s Rule
3. Pauli’s Exclusion Principle
13
1.Aufbau (building up) PrincipleElectrons enter the orbitals in order of ascending energy.
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s p d f g h i
Letters read across
1
2
3
4
5
6
7
Numbers read downwards
2
3 3
4 4 4
5 5 5 5
6 6 6 6 6
7 7 7 7 7 7
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s p d f g h i1
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
1s,
16
s p d f g h i1
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
1s, 2s,
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s p d f g h i1
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
1s, 2s, 2p, 3s,
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s p d f g h i1
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
1s, 2s, 2p, 3s, 3p, 4s,
19
s p d f g h i1
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
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s
20
s p d f g h i1
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
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s
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s p d f g h i1
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
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s
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s p d f g h i1
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
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s
23
Building up of electronic Building up of electronic configurationsconfigurations
5.1 Relative energies of orbitals (SB p.106)
24
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
25
3. Pauli’s exclusion principle : -
Electrons occupying the same orbital must have opposite spins.
W. Pauli , Nobel prize laureate in Physics, 1945
Check Point 5-1Check Point 5-1
26
5.5.22 Electronic Electronic ConfigurationConfigurations of Elementss of Elements
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Ways to Express Electronic ConfigurationsWays 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, 2py
2, 2pz2, 3s1
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K 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Fe 1s2, 2s2, 2p6, 3s2, 3p6, 3d6, 4s2
Q.17(a)
Q.17(b)
orbitals of the same quantum shell are placed together
29
2. Using a noble gas ‘core’
Na [Ne] 3s1
Ca [Ar] 4s2
30
Si [Ne] 3s2, 3p2
V [Ar] 3d3, 4s2
Q.18(a)
Q.18(b)
outermost shell
outermost shell
31
3. Electrons – in – Boxes representation
N1s 2s 2p
(1) All boxes should be labelled
(2) Boxes of the same energies are put together.
2px 2py 2pz
32
Q.19
Hund’s rule is violatedHund’s rule is violated
Pauli’s exclusion principle is violated
33
Q.20(a)
1s 2s 2p 3p3s
Phosphorus
34
Q.20(b)
Chromium1s 2s 3s2p 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.
35
Q.20(b)Chromium
[Ar] 3d4, 4s2 [Ar] 3d5, 4s1 + energy
+ energy
1s 2s 3s2p 3p 3d 4s
1s 2s 3s2p 3p 3d 4s
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.
1s 2s 3s2p 3p 3d 4s
37
[Ar] 3d9, 4s2 [Ar] 3d10, 4s1 + energyQ.20(c)
Copper
+ energy
1s 2s 3s2p 3p 4s
1s 2s 3s2p 3p 4s
3d
3d
38
Silicon3p3s
[Ne]
Empty orbital(s) in a partially filled subshell should be shown
39
Silicon3p3s
[Ne]
3p3s
[Ne]
+ energy
Energy difference : 3p – 3s > 3d – 4s
40
21(a)(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
21(b) Ar
41
S2 , Cl , Ar , K+ , Ca2+
Same electronic configurations
IsoelectronicQ.22
42
Represented by ‘electrons-in-boxes’ Represented by ‘electrons-in-boxes’ diagramsdiagrams
5.2 Electronic configurations of elements (SB p.110)
43
5.2 Electronic configurations of elements (SB p.110) Check Point 5-2Check Point 5-2
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http://www.chemcollective.org/applets/pertable.php
Building up of electronic Building up of electronic configurationsconfigurations
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 : 地、水、火、風Quintessence (The fifth element)
、空
47
Seven Planetary Elements of Alchemists
48
Sun Gold
Moon Silver Venus
Copper
Mars Iron
Jupiter TinMercury Mercury Saturn
Lead
49
Other Alchemical Elements
Sb
PtBi
As
SP
50
Law of Triads (Dobereiner, 1829)
Element Molar mass(g/mol)
Density(g/cm³)
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 137.327 3.594
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
Group 1A
Group 2A
Group 3A
Group 4A
Group 5A
Group 6A
Group 7A
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
‘missing elements’ predicted by Mendeleev
1. Ekaboron (atomic mass = 44)
Scandium (44.96)
2. Ekaaluminium (68)
Gallium (69.3)
First Periodic Table (Mendeleev, 1869)
54
‘missing elements’ predicted by Mendeleev
First Periodic Table (Mendeleev, 1869)
3. Ekamanganese (100)
Technetium (98)
4. Ekasilicon (72)
Germanium (72.59)
55
7 groups or 8 groups ?
56
Discovery of the Noble Gases
Lord Rayleigh William RamsayNobel Laureate
in Physics, 1904Nobel Laureate in Chemistry, 1904
57
Air - (O2, CO2, H2O)
N2
Density ( g / dm3)
1.2572
NH3 N2 1.2508
% error 0.5% ???
decompose
1894
58
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
59
Group 1A
Group 2A
Group 3A
Group 4A
Group 5A
Group 6A
Group 7A
Group 0
Li Be B C N O F Ne
Na Mg Al Si P S Cl Ar
Rn discovered in 1900 by F.E. Dorn
Helium discovered in 1895
Ne, Kr, Xe discovered in 1898All by Ramsay
Po, Ra discovered in 1898 by Pierre & Marie Curie
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Congratulations !
Nobel Laureate in Chemistry,
2010
61
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
Modern Periodic Table
62
Transuranium Elements
Discovered by McMillan and Seaborg
99989694969594
Pu93
Np92
U
63
Transuranium Elements
Discovered by McMillan and Seaborg
Nobel Laureates in Chemistry, 1951
From University of California, Berkeley,
United States of America
92
U93
Np94
Pu
64
92
U93 94 95 96 94 96 98 99
Uranium, discovered in 1789,
was considered the heaviest elements
65
92
U
Named after Uranus ( 天王星 )
Discovered in 1781
Was Considered the Farthest Planet from The Earth in the Solar System
66
92
U93
Np94
Pu
Transuranium Elements
Neptune( 海王星 ) : The Next Planet out from Uranus
Neptunium : Discovered in 1940 by McMillan
67
92
U93
Np94
Pu
Transuranium Elements
Pluto( 冥王星 ) : Was considered the next ‘Planet’ out from Neptune
Plutonium : Discovered in 1941 by McMillan & Seaborg
68
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureates in Chemistry, 1951
University of California, Berkeley,
United States of America
Americium (1944)
69
It was named americium because it is just below europium in the Periodic Table.
70
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureate in Physics, 1903
Nobel Laureate in Chemistry, 1911
Curium Marie Curie
71
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureates in Chemistry, 1951
University of California, Berkeley,
United States of America
Berkelium (1949)
72
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureates in Chemistry, 1951
University of California, Berkeley,
United States of America
Californium (1950)
73
Transuranium Elements
McMillan and Seaborg
Nobel Laureates in Chemistry, 1951
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
74
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureate in Physics, 1921
Einsteinium (1952 by Albert Ghiorso)
Albert Einstein
75
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureate in Physics, 1938
Developer of the first nuclear reactor, 1942
Fermium (1952 by Albert Ghiorso)Enrico Fermi
76
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
General Consultant of the Manhattan Project
Hiroshima – little boy
Nagasaki – fat man
77
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Discovery of Periodicity 1869
Mendelevium (1955 by Albert Ghiorso)
Mendeléev
78
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Inventor of Dynamite, 1867
The Man Behind the Nobel Prize
Nobelium (1958 by Albert Ghiorso)
Alfred Nobel
79
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Lawrencium (1961 by Albert Ghiorso)
@ Lawrence Radiation Laboratory University of California, Berkeley,
80
95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
Transuranium Elements
Nobel Laureate in Physics, 1939
Ernest Orlando Lawrence
University of California, Berkeley,
Developer of cyclotron
81
1964 – 1996 ADTeams 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 System0 = nil
1 = un
2 = bi
3 = tri
4 = quad
5 = pent
6 = hex
7 = sept8 = oct9 = enn
ium ium
111 = unununium (Uuu) = Roentgenium (Rg)
84
111 = unununium (Uuu)112 = ununbium (Uub)113 = ununtrium (Uut)
114 = ununquadium (Uuq)115 = ununpentium (Uup)
116 = ununhexium (Uuh)
d-block
p-block
6B
85
The Periodic TableThe Periodic Table
5.3 The Periodic Table (SB p.112)
86
d-block
p-block
f-block
s-block
5.3 The Periodic Table (SB p.112)
s-block & p-block elements are called representative elements
87
f-block elements are called inner-transition elements
Rare Earth Metals( 稀土金屬 )
88
中東有石油中國有稀土
鄧小平
89
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 )
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 – 1n 4
n – 2n 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
3
4
5
6
7
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)
VIIIB Fe [Ar] 3d6, 4s2
Co [Ar] 3d7, 4s2
Ni [Ar] 3d8, 4s2
Not true for VIIIB elements
97
5.3 The Periodic Table (SB p.112)
Check Point 5-3Check Point 5-3
98
The Song of Elements by Tom Lehrer
The Song of Elements – on YouTube
Visual Elements Periodic Table
99
Periodicity as illustrated by
(i) Variation in atomic radius with atomic number
(ii) Variation in ionization enthalpy with atomic number
100
5.5.66 Atomic Size Atomic Size
of Elementsof 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
Q: Explain why the atomic radius decreases across a period.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.
5.6 Atomic size of elements (p. 122)
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
Two electrons in the inner shell
nucleus
2s electron of Li outside the inner shell
108
The outer 2s electron is repelled/shielded/screened by the inner 1s electrons from the nucleus
+3
Li
109
Li
+1
The nuclear charge experienced by the 2s electron is +1
+3
Li
110
+4
Be
The inner 1s electrons shield the outer electrons almost completely
Be
+2
111
Be
+2
The two electrons in the same shell (2s) shield each other less poorly.
Be
+1.5
Zeff 1.5
112
113
Atomic radii down a group
114
Q: Explain why the atomic radius increases down a group.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.
• 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.
5.6 Atomic size of elements (p. 122)
115
Sharp in atomic radius when a new Period begins
116
Q: Explain why there is sharp in atomic radius when a new Period beginsQ: 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.
5.4 Ionization enthalpies of elements (SB p.117)
117
Q: Explain why there is sharp in atomic radius when a new Period beginsQ: 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.
5.4 Ionization enthalpies of elements (SB p.117)
Check Point 5-6Check Point 5-6
118
5.5.44 Ionization Ionization Enthalpies of Enthalpies of
ElementsElements
119
Across a Period, there is a general in I.E. leading to a maximum with a noble gas.
120
Effective nuclear charge from left to right across the Period
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
The first ionization enthalpies generally decrease down a group and increase across a period
5.4 Ionization enthalpies of elements (SB p.116)
124
125
Q: Explain why there is sharp in IE when a new Period begins
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.
5.4 Ionization enthalpies of elements (SB p.117)
He 1s2 (duplet)
Ne 2s2, 2p6 (octet) Fully-filled
shells
Ar 3s2, 3p6 (octet) Fully-filled subshell
126
Q: Explain why there is sharp in I.E. when a new Period beginsQ: 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.
5.4 Ionization enthalpies of elements (SB p.117)
127
Irregularities : -
128
Q: Explain why there is a trough at Boron(B) in Period 2.Q: Explain why there is a trough at Boron(B) in Period 2.
• Be : 1s2, 2s2
B : 1s2, 2s2, 2p1• Be : 1s2, 2s2
B : 1s2, 2s2, 2p1
5.4 Ionization enthalpies of elements (SB p.117)
129
More diffused
More penetrating
1s
130
In multi-electron systems,
penetrating power : -
s > p > d > f
131
22 4 r
3d electrons are more diffused (less penetrating)
3d electrons are more shielded by 1s electrons
132
Q: Explain why there is a trough at Boron(B) in Period 2.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.
• 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.
5.4 Ionization enthalpies of elements (SB p.117)
133
Irregularities : -
134
Q: Explain why there is a trough at Oxygen(O) in Period 2.Q: Explain why there is a trough at Oxygen(O) in Period 2.
• e.c. of N : 1s2, 2s2, 2px1, 2py
1, 2pz1
e.c. of O : 1s2, 2s2, 2px2, 2py
1, 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.
• e.c. of N : 1s2, 2s2, 2px1, 2py
1, 2pz1
e.c. of O : 1s2, 2s2, 2px2, 2py
1, 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.
5.4 Ionization enthalpies of elements (SB p.117)
135
Q: Explain why there is a trough at Oxygen(O) in Period 2.Q: Explain why there is a trough at Oxygen(O) in Period 2.
e.c. of N : 1s2, 2s2, 2px1, 2py
1, 2pz1
e.c. of O : 1s2, 2s2, 2px2, 2py
1, 2pz1
Alternately, the removal of a 2p electronfrom O results in a stable half-filled 2psubshell.
e.c. of N : 1s2, 2s2, 2px1, 2py
1, 2pz1
e.c. of O : 1s2, 2s2, 2px2, 2py
1, 2pz1
Alternately, the removal of a 2p electronfrom O results in a stable half-filled 2psubshell.
5.4 Ionization enthalpies of elements (SB p.117)
136
5.5.55Variation of Variation of
Successive Successive Ionization Ionization
Enthalpies with Enthalpies with Atomic NumbersAtomic 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. > 2nd I.E. > 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. Be+ = Li (2, 1)
B+ = Be (2, 2)
C+ = B (2, 3)
Each represents a pair of isoelectronic species
139
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 ?
Invert relationship between atomic radius and first I.E.
140
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.
141
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.
142
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
143
The END
144
Write the electronic configurations and draw “electrons-in –boxes” diagrams for
(a) nitrogen; and
(b) sodium.
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Answer
5.1 Relative energies of orbitals (SB p.108)
(a) Nitrogen: 1s22s22p3
(b) Sodium: 1s22s22p63s1
145
Give the electronic configuration by notations and “electrons-in-boxes” diagrams in the abbreviated form for the following elements.
(a) silicon; and
(b) copper.
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Answer
5.2 Electronic configurations of elements (SB p.110)
(a) Silicon: [Ne]3s23p3
(b) Copper: [Ar]3d104s1
146
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?
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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.
Answer
5.3 The Periodic Table (SB p.113)
147
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
148
5.3 The Periodic Table (SB p.114)
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149
(a)Give four main factors that affect the magnitude of ionization enthalpy of an atom.
Answer
5.4 Ionization enthalpies of elements (SB p.118)
(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.
150
(b)Explain why Group 0 elements have extra high first ionization enthalpies and their decreasing trend down the group.
Answer
5.4 Ionization enthalpies of elements (SB p.118)
(b) 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
(c)Predict the trend of the first ionization enthalpies of the transition elements.
Answer
5.4 Ionization enthalpies of elements (SB p.118)
(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.
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152
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-in- boxes” diagram. Answer(a) (i) 1s22s22p2
(ii)
153
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
(b)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. Answer
1st 2nd 3rd 4th 5th 6th
I.E. (kJ
mol-1)
1090 2350 4610 6220 37800
47000
154
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
(b) (i)
155
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
(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.
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156
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(a) Give the “electrons-in-boxes” diagram of 26Fe.
(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+.
Answer
(a) Fe :
(b) Fe2+ :
Fe3+ :
157
(c) Which ion is more stable, Fe2+ or Fe3+? Explain briefly.
Answer
(c) Fe3+ ion is more stable because the 3d orbital is exactly half-filled which gives the electronic configuration extra stability.
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
158
(d)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.
Answer
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
1st 2nd 3rd 4th 5th 6th
I.E. (kJ
mol-1)
762 1560 2960 5400 7620 10100
159
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
(d) (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)
(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 su
ccessive ionization enthalpies of iron.
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161
Explain the following:
(a) The atomic radius decreases across the period from Li to Ne.
5.6 Atomic size of elements (p. 123)
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
Explain the following:
(b) The atomic radius increases down Group I metals.
5.6 Atomic size of elements (p. 123)
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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.