The Electronic Structure of Atoms

1
The Electronic Structure of Atoms
4.1 The Electromagnetic Spectrum 4.2 Deduction of
Electronic Structure from Ionization
Enthalpies 4.3 The Wave-mechanical Model of the
Atom 4.4 Atomic Orbitals
2
The electronic structure of atoms
Chapter 4 The electronic structure of atoms (SB
p.80)

• Two sources of evidence -
• Study of atomic emission spectra of the
  elements
• Study of ionization enthalpies of the elements

3
Atomic Emission Spectra??????
Spectra ? plural of Spectrum Arises from light
emitted from individual atoms
4
The electromagnetic spectrum
Wavelength (m) ?
? Frequency (Hz / s?1)
Speed of light (ms?1) Frequency ? Wavelength
c ?? ? 3?108 ms?1
5
The electromagnetic spectrum
red
Wavelength (m) ?
violet
? Frequency (Hz / s?1)
6
The electromagnetic spectrum
red
Wavelength (m) ?
violet
? Frequency (Hz / s?1)
? Increasing energy
7
The electromagnetic spectrum
red
Wavelength (m) ?
violet
? Frequency (Hz / s?1)
? Increasing energy
8
The electromagnetic spectrum
red
Wavelength (m) ?
violet
? Frequency (Hz / s?1)
? Increasing energy
9
The electromagnetic spectrum
red
Wavelength (m) ?
violet
? Frequency (Hz / s?1)
Decreasing energy ?
10
The electromagnetic spectrum
red
Wavelength (m) ?
violet
? Frequency (Hz / s?1)
Decreasing energy ?
11
Types of Emission Spectra
4.1 The electromagnetic spectrum (SB p.82)

• Continuous spectra
• E.g. Spectra from tungsten filament and
  sunlight
• Line Spectra
• E.g. Spectra from excited samples in
  discharge tubes

12
Continuous spectrum of white light
4.1 The electromagnetic spectrum (SB p.82)
Fig.4-5(a)
13
Line spectrum of hydrogen
4.1 The electromagnetic spectrum (SB p.83)
Fig.4-5(b)
14
How Do Atoms Emit Light ?
Hydrogen atom in ground state means its electron
has the lowest energy
15
Atoms in excited state
16
Not Stable
Atom returns to ground state
17
Atom returns to ground state
E h?
Plancks equation
18
Atom returns to ground state
E h?
Planck Nobel laureate Physics, 1918
19
E2
E1
Atom returns to ground state
E h?
E energy of the emitted light E2 E1
20
E2
E1
Atom returns to ground state
E h?
? Frequency of the emitted light
21
E2
E1
Atom returns to ground state
E h?
h (Plancks constant) 6.63 ? 10?34 Js
22
E2
E1
Atom returns to ground state
E h?
h 6.63 ? 10?34 Js
Energy cannot be absorbed or emitted by an atom
in any arbitrary amount.
23
E2
E1
Atom returns to ground state
E h?
h 6.63 ? 10?34 Js
Energy can only be absorbed or emitted by an atom
in multiples of 6.63?10?34 J.
24
Characteristic Features of the Hydrogen Emission
Line Spectrum
4.1 The electromagnetic spectrum (SB p.84)
1. The visible region The Balmer Series
25
Q.1
3.03?10?19 J
26
energy of one photon emitted
27
Q.2
(a) The spectral lines come closer at higher
frequency and eventually merge into a
continuum(???) (b) n ? ? n 2
28
Q.2
(c) The electron has been removed from the
atom. I.e. the atom has been ionized.
H(g) ? H(g) e?
29
corresponds to the last spectral line of the
Balmer series
30
The Complete Hydrogen Emission Spectrum
4.1 The electromagnetic spectrum (SB p.83)
UV
Visible
IR
Let's Think 1
31
Q.3
(a) The spectral lines in each series get
closer at higher frequency. (b) Since the energy
levels converge at higher level, the spectral
lines also converge at higher frequency.
32
Rydberg Equation
Relates wavelength of the emitted light of
hydrogen atom with the electron transition
33
number of waves in a unit length
? 100 waves in 1 meter
34
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35
Electron transition b ? a, a, b are integers
and b gt a a represents the lower energy level to
which the electron is dropping back b represents
the higher energy levels from which the electron
is dropping back
36
Balmer series, a 2 b 3, 4, 5,
37
Lyman series, a 1 b 2, 3, 4,
2 3 4 5 6 ?
38
Paschen series, a 3 b 4, 5, 6,
4 5 6 7 ?
39
RH is the Rydberg constant
40
Q.4
(106 m?1)
41
RH 1.096?107 m?1
42
Interpretation of the atomic hydrogen spectrum
4.1 The electromagnetic spectrum (SB p.84)
Discrete spectral lines ? energy possessed by
electrons within hydrogen atoms cannot be in
any arbitrary quantities but only in specified
amounts called quanta.
43
Interpretation of the atomic hydrogen spectrum
4.1 The electromagnetic spectrum (SB p.84)
Only certain energy levels are allowed for the
electron in a hydrogen atom. The energy of the
electron in a hydrogen atom is quantized.
44
Interpretation of the atomic hydrogen spectrum
4.1 The electromagnetic spectrum (SB p.84)
45
Interpretation of the atomic hydrogen spectrum
Bohrs Atomic Model of Hydrogen
Nobel Prize Laureate in Physics, 1922
46
Interpretation of the atomic hydrogen spectrum
Nobel Prize Laureate in Physics, 1922
for his services in the investigation of the
structure of atoms and of the radiation emanating
from them
47
Interpretation of the atomic hydrogen spectrum
Bohrs model of H atom
48
Interpretation of the atomic hydrogen spectrum

• The electron can only move around the nucleus of
  a hydrogen atom in certain circular orbits with
  fixed radii.
• Each allowed orbit is assigned an integer, n,
  known as the principal quantum number.

49
Interpretation of the atomic hydrogen spectrum
2. Different orbits have different energy
levels. An orbit with higher energy is further
away from the nucleus.
50
Interpretation of the atomic hydrogen spectrum

• Spectral lines arise from electron transitions
  from higher orbits to lower orbits.

51
Interpretation of the atomic hydrogen spectrum
For the electron transition E2 ? E1,
?E E2 E1 h?
52
Interpretation of the atomic hydrogen spectrum

• In a sample containing numerous excited H atoms,
• different H atoms may undergo different kinds
  of electron transitions to give a complete
  emission spectrum.

53
Wavelength Determined by Experiment (nm)
Wavelength Predicted by Bohr (nm)
Electronic Transition
2 ? 1
3 ? 1
4 ? 1
3 ? 2
4 ? 2
5 ? 2
4 ? 3
54
Interpretation of the atomic hydrogen spectrum
5. The theory failed when applied to elements
other than hydrogen (multi-electron systems)
55
Illustrating Bohrs Theory
First line of Lyman series, n 2 ? n 1
By Plancks equation,
56
By Rydbergs equation,
57
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58
All electric potential energies have negative
signs except E? E1 lt E2 lt E3 lt E4 lt E5 lt lt
E? 0
59
Balmer series, Transitions from higher levels to
n 2
Visible region
60
Lyman series, Transitions from higher levels to n
1
More energy released ? Ultraviolet region
61
Paschen series, Transitions from higher levels to
n 3
Less energy released ? Infra-red region
62
4.1 The electromagnetic spectrum (SB p.87)
?E E2 E1 h?
63
Q.5 n 100 ? n 99
Energy of the light emitted is extremely small.
64
The electromagnetic spectrum
Wavelength (m) ?
? Frequency (Hz / s?1)
65
Q.5 n 100 ? n 99
Energy of the light emitted is extremely small.
? Microwave region
66
Q.6
67
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68
45.7?1013 Hz
69
Energy of the second line
70
61.7?1013 Hz
71
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72
69.1?1013 Hz
73
Convergence Limits and Ionization Enthalpies
Ionization enthalpy is the energy needed to
remove one mole of electrons from one mole of
gaseous atoms in ground state to give one moles
of gaseous ions (n ?)
X(g) ? X(g) e?
Units kJ mol?1
74
Convergence Limits and Ionization Enthalpies
The frequency at which the spectral lines of a
series merge.
75
Q.7
H(g) ? H(g) e?
Ionization enthalpy
76
Electron transition for ionization of atom
Energy Level of e? to be removed in ground state
of atom
Atom
H
He
Li
Na
77
Q.9
Ionization enthalpy of helium
(6.626?10?34 Js)(5.29?1015 s?1)(6.02?1023
mol?1) 2110 KJ mol?1 gt Ionization enthaply of
hydrogen
78
He
H
Relative positions of energy levels depend on the
nuclear charge of the atom.
E2
E2
E1
E1
Ionization enthalpy He gt H
79
The uniqueness of atomic emission spectra
4.1 The electromagnetic spectrum (SB p.89)
No two elements have identical atomic spectra
80
The uniqueness of atomic emission spectra
4.1 The electromagnetic spectrum (SB p.89)
atomic spectra can be used to identify unknown
elements.
81
Flame Test
82
Flame Test
83
Q.10
4.42?10?19 J
84
Emission vs Absorption
Bright lines against a dark background
Dark lines against a bright background
85
Absorption spectra are used to determine the
distances and chemical compositions of the
invisible clouds.
86
The Doppler effect
The frequency of sound waves from a moving object
(a) increases when the object moves towards the
observer. (b) decreases when the object moves
away from the observer.
87
Redshift and the Doppler effect

• Frequency of light waves emanating from a moving
  object
• decreases when the light source moves away from
  the observer.
• wavelength increases
• spectral lines shift to the red side with
  longer wavelength
• redshift

88
Atomic Absorption Spectra
Moving at higher speed
Moving at lower speed
89
Redshift
Left - Absorption spectrum from
sunlight. Right - Absorption
spectrum from a supercluster of distant galaxies
90
Atomic Absorption Spectra
Only atomic emission spectrum of hydrogen is
required in A-level syllabus !!!
91
Deduction of Electronic Structure from Ionization
Enthalpies
92
Evidence of the Existence of Shells
For multi-electron systems, Two questions need to
be answered
93
Evidence of the Existence of Shells

• How many electrons are allowed to occupy each
  electron shells ?

Existence of Shells
2. How are the electrons arranged in each
electron shell ?
Existence of Subshells
94
Ionization enthalpy
Ionization enthalpy is the energy needed to
remove one mole of electrons from one mole of
gaseous atoms in ground state to give one moles
of gaseous ions (n ?)
X(g) ? X(g) e?
95
Evidence of the Existence of Shells
Successive ionization enthalpies Q. 11
96
Q.11(a)
Be(g) ? Be(g) e? IE1
Be(g) ? Be2(g) e? IE2
Be2(g) ? Be3(g) e? IE3
Be3(g) ? Be4(g) e? IE4
97
Q.11(b)
IE1 lt IE2 lt IE3 lt IE4
Positive ions with higher charges attract
electrons more strongly. Thus, more energy is
needed to remove an electron from positive ions
with higher charges.
98
Q.11(c)
99
Q.11(c)
(i) The first two electrons are relatively easy
to be removed. ? they experience less
attraction from the nucleus, ?
they are further away from the
nucleus and
occupy the n 2 electron shell.
100
Q.11(c)
(ii) The last two electrons are very difficult
to be removed. ? they experience stronger
attraction from the nucleus, ? they are
close to the nucleus and
occupy the n 1 electron shell.
101
Electron Diagram of Beryllium
102
Energy Level Diagram of Beryllium
103
IE1 E? - E2
104
IE1
IE2
IE1 lt IE2
105
IE1
IE2
IE3
IE1 lt IE2 ltlt IE3
106
Be
IE1
IE2
IE3
IE4
IE1 lt IE2 ltlt IE3 lt IE4
107
The Concept of Spin(??)
Spin is the angular momentum intrinsic to a
body. E.g. Earths spin is the angular momentum
associated with Earths rotation about its own
axis.
??
108
On the other hand, orbital angular momentum of
the Earth is associated with its annual motion
around the Sun (??)
109
Subatomic particles like protons and electrons
possess spin properties. i.e. they have spin
angular momentum. But their spins cannot be
associated with rotation since they display both
particle-like and wave-like behaviours.
110
Paired electrons in an energy level should have
opposite spins. Electrons with opposite spins are
represented by arrows in opposite directions.
Q.12
111
Q.12
4 groups of electrons
112
n 1
Which group of electrons is in the first shell ?
Q.12
n 4
113
Q.12
2, 8, 8, 2
114
Q.12
115
Evidence of the Existence of Shells
4.2 Deduction of electronic structure from
ionization enthalpies (p.91)
2, 8, 1
116
Na
117
Variation of IE1 with Atomic Number
Evidence for Subshell
118
Only patterns across periods are discussed
Refer to pp.27-29 for further discussion
119
1. A general ? in IE1 with atomic number across
Periods 2 and 3.
13.(a)
120
2 3 - 3
13.(a)
121

• IE1 value Group 2 gt Group 3

13.(a)
3. IE1 value Group 5 gt Group 6
122

• Peaks appear at Groups 2 5

13.(a)
3. Troughs appear at Groups 3 6
123
On moving across a period from left to right,
1. the nuclear charge of the atoms ? (from 3 to
10 or 11 to 18)
13.(b)
2. electrons are being removed from the same shell
? e?s removed experience stronger attraction from
the nucleus.
124
IE1 B(2,3) lt Be(2,2) The electron removed from
B occupies a subshell of higher energy within the
n 2 quantum shell
13.(b)
125
13.(b)
IE2s(Be) gt IE2p(B) IE1 Be gt B
1s
126
IE1 Al(2,8,3) lt Mg(2,8,2) The electron removed
from Al occupies a subshell of higher energy
within the n 3 quantum shell
13.(b)
127
13.(b)
IE3s(Mg) gt IE3p(Al) IE1 Mg gt Al
2p
128
n ?
2p
Mg(12)
Al(13)
129
P(2,8,5) gt S(2,8,6)
IE1 N(2,5) gt O(2,6)
It is more difficult to remove an electron from a
half-filled p subshell
13.(b)
130
2p
2s
131
2p
2s
The three electrons in the half-filled 2p
subshell occupy three different orbitals (2px ,
2py , 2pz).
? repulsion between electrons is minimized.
132
3p
3s
The three electrons in the half-filled 3p
subshell occupy three different orbitals (3px ,
3py , 3pz). ? repulsion between electrons is
minimized.
133
The removal of an electron from O or S results in
a half-filled p subshell with extra
stability.(p.28, part 3)
Misleading !!!
134
O(2,6) ? O(2,5)
2p
Is the 2p energy level of O lower or higher than
that of O ?
Electrons in O experience a stronger attractive
force from the nucleus.
Not because of the half-filled 2p subshell
135
O(2,6) ? O(2,5)
2p
As a whole, O is always less stable than O.
It is because O(g) has one more electron than
O(g) and this extra electron has a negative
potential energy.
136
Conclusions -
(a) Each electron in an atom is described by a
set of four quantum numbers. (b) No two electrons
in the same atom can have the same set of
quantum numbers.
The quantum numbers can be obtained by solving
the Schrodinger equation (p.16).
137
Quantum Numbers

• Principal quantum number (n)
• n 1, 2, 3,
• related to the size and energy of the principal
  quantum shell.
• E.g. n 1 shell has the smallest size and
  electrons in it possess the lowest energy

138
Quantum Numbers
related to the shape of the subshells.
139
Quantum Numbers
140
Each principal quantum shell can have one or more
subshells depending on the value of n.
If n 1,
0 to (1-1)
? 0
No. of subshell 1
name of subshell 1s
141
Each principal quantum shell can have one or more
subshells depending on the value of n.
If n 2,
0 to (2-1) ? 0, 1
No. of subshells 2
names of subshells 2s, 2p
142
Each principal quantum shell can have one or more
subshells depending on the value of n.
If n 3,
0 to (3-1) ? 0, 1, 2
No. of subshells 3
names of subshells 3s, 3p, 3d
143
Each principal quantum shell can have one or more
subshells depending on the value of n.
If n 4,
0 to (4-1) ? 0, 1, 2, 3
No. of subshells 4
names of subshells 4s, 4p, 4d, 4f
144
Quantum Numbers
3. Magnetic quantum number (m)
related to the spatial orientation of the
orbitals in a magnetic field.
145
Possible range of m ?0 to 0 ? 0 No. of orbital
1 Name of orbital s
146
Possible range of m ?1 to 1 ? ?1 ,0,1 No. of
orbitals 3 Names of orbitals px, py, pz
147
Possible range of m
?2 to 2 ? ?2, ?1, 0, 1, 2 No. of orbitals
5 Names of orbitals
148
Possible range of m
?3 to 3 ? ?3, ?2, ?1, 0, 1, 2, 3 No. of
orbitals 7 Names of orbitals Not required in
AL
149
Total no. of orbitals in a subshell
Energy of subshells - s lt p lt d lt f
150

• Spin quantum number (ms)

They describe the spin property of the electron,
either clockwise, or anti-clockwise
151

• Spin quantum number (ms)

Each orbital can accommodate a maximum of two
electrons with opposite spins
152
Q.14(a)
Total no. of electrons
Total no. of orbitals
Subshells
Principal quantum shell
2
1(1s)
1s
n 1
8
n 2
134
18
n 3
1359
32
n 4
135716
153
14(b)
The total number of electrons in a principal
quantum shell 2n2
154
Q.15(a)
The two electrons of helium are in the 1s
orbital of the
1s subshell of the
first principal quantum shell.
155
Q.15(b)
There are only 2 electrons in the 3rd quantum
shell. In the ground state, these two electrons
should occupy the 3s subshell since electrons in
it have the lowest energy.
156
Q.15(b)
The two outermost electrons of magnesium are in
the
3s orbital of the
3s subshell of the
third principal quantum shell.
157
The Wave-mechanical Model of the Atom
158
Models of the Atoms

• Plum-pudding model by J.J. Thomson (1899)
• Planetary(orbit)model by Niels Bohr (1913)
• Orbital model by E. Schrodinger (1920s)

159
Electrons display both particle nature and wave
nature. Particle nature mass, momemtum, Wave
nature frequency, wavelength, diffraction,
160
Wave as particles photons
momentum of a photon
161
Evidence photoelectric effect by
Albert Einstein (1905) Nobel Prize Laureate in
Physics, 1921
162
Particles as waves
L. De Broglie (1924) Nobel Prize Laureate in
Physics, 1929
163
(3)
(4)
Any particle (not only photon) in motion (with a
momentum, mv) is associated with a wavelength
164
Q.16
Electron
165
Q.16
Helium atom
166
Q.16
100m world record holder
167
Wave nature of electrons (1927)
4.3 The Wave-mechanical model of the atom (p.95)
Evidence
? of electron
? inter-atomic spacing in
metallic crystals (10?10m)
168
For very massive particles, The wavelength
associated with them (10?37m) are much smaller
than the dimensions of any physical system. Wave
properties cannot be observed.
169
Standing waves in a cavity
Standing waves only have certain allowable modes
of vibration Similarly, electrons as waves only
have certain allowable energies.
170
The uncertainty principle
Heinsenberg Nobel Prize Laureate in Physics, 1932
171
The uncertainty principle
It is impossible to simultaneously determine the
exact position and the exact momentum of an
electron.
172
The uncertainty principle
The momentum of electron would change greatly
after collision
173
The uncertainty principle
The position of electron cannot be located
accurately
174
The uncertainty principle
No problem in macroscopic world !
175
Implications - The concept of well-defined
orbits in Bohrs model has to be abandoned. We
can only consider the probability of finding an
electron of a certain energy and momentum
within a given space.
176
Schrodinger Equation
Nobel Prize Laureate in Physics, 1933
de Broglie electrons as waves ? Use wave
functions (?) to describe electrons
177
Schrodinger Equation
Nobel Prize Laureate in Physics, 1933
Heisenberg Uncertainty principle ?
Probability (?2) of finding electron at a
certain position lt 1.
178
Schrodinger Equation
? wave function m mass of electron h
Plancks constant E Total energy of electron V
Potential energy of electron
179
Schrodinger Equation
The equation can only be solved for certain ?i
and Ei
180
Schrodinger Equation
The wave function of an 1s electron is
Z nuclear charge (Z 1 for Hydrogen) a0 Bohr
radius 0.529Å (1Å 10?10m) r distance of
electron from the nucleus
181
Schrodinger Equation
The allowed energies of H atom are given by
n is the principal quantum number, All other
terms in the expression are constants
182
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183
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184
?2is the probability of finding an electron at a
particular point in space. (electron density)
4.4 Atomic orbitals (p.98)
Probability never becomes zero ? There is no
limit to the size of an atom
185
relative probability of finding the electron
at the nucleus
contour diagram
186
In practice, a boundary surface
is chosen such that within which there is a high
probability (e.g. 90) of finding the electron.
187
The electron spends 90 of time within the
boundary surface
188
A 3-dimensional time exposure diagram. The
density of the dots represents the probability of
finding the electron at that position.
189
The 3-dimensional region within which there is a
high probability of finding an electron in an
atom is called an atomic orbital. Each atomic
orbital is represented by a specific wave
function(?). The wave function of a specific
atomic orbital describes the behaviour of the
electron in the orbital.
190
dr is infinitesimally small
Total probability of finding the electron within
the shell of thickness dr ?24?r2dr
?2 is the probability of finding the electron per
unit volume
191
?2? as r ? , 4?r2 ? as r ? ? a maximum at 0.529
Å
Orbital Model
192
Bohrs Orbit Model
193
Total probability of finding the electron within
the shell of thickness dr ?24?r2dr
The sum of the probabilities of finding the
electron within all shells 1
194
The total area bounded by the curve and the
x-axis 1
195
s Orbitals
4.4 Atomic orbitals (p.98)
196
There is no chance of finding the
electron on the nodal surface.
197
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198
How can the electron move between A B ?
199
? can be considered as the amplitude of the wave.
?2 is always ? 0
200
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201
p Orbitals
4.4 Atomic orbitals (p.100)
Two lobes along an axis
202
For each 2p orbital, there is a nodal plane on
which the probability of finding the electron is
zero.
4.4 Atomic orbitals (p.100)
yz plane
xz plane
xy plane
203
d Orbitals
4.4 Atomic orbitals (p.101)
Four lobes between two axes
Four lobes along two axes
Two lobes one belt
204
Grand Orbital Table
http//www.orbitals.com/orb/orbtable.htmtable1
205
The END
206
4.1 The electromagnetic spectrum (SB p.82)
Let's Think 1
Some insects, such as bees, can see light of
shorter wavelengths than humans can. What kind of
radiation do you think a bee sees?
Answer
Ultraviolet radiation
Back
207
4.1 The electromagnetic spectrum (SB p.87)
Let's Think 2
What does the convergence limit in the Balmer
series correspond to?
Answer
The convergence limit in the Balmer series
corresponds to the electronic transition from n
? to n 2.
Back
208
4.1 The electromagnetic spectrum (SB p.88)
Example 4-1A
Given the frequency of the convergence limit of
the Lyman series of hydrogen, find the ionization
enthalpy of hydrogen. Frequency of the
convergence limit 3.29 ? 1015 Hz Planck
constant 6.626 ? 10-34 J s Avogadro constant
6.02 ? 1023 mol-1
Answer
209
4.1 The electromagnetic spectrum (SB p.88)
Back
Example 4-1A
For one hydrogen atom, E h? 6.626 ? 10-34
J s ? 3.29 ? 1015 s-1 2.18 ? 10-18 J For
one mole of hydrogen atoms, E 2.18 ? 10-18 J ?
6.02 ? 1023 mol-1 1312360 J mol-1
1312 kJ mol-1 The ionization enthalpy of hydrogen
is 1312 kJ mol-1.
210
4.1 The electromagnetic spectrum (SB p.88)
Example 4-1B
The emission spectrum of atomic sodium is
studied. The wavelength of the convergence limit
corresponding to the ionization of a sodium atom
is found. Based on this wavelength, find the
ionization enthalpy of sodium. Wavelength of the
convergence limit 242 nm Planck constant
6.626 ? 10-34 J s Avogadro constant 6.02 ? 1023
mol-1 Speed of light 3 ? 108 m s-1 1 nm 10-9 m
Answer
211
4.1 The electromagnetic spectrum (SB p.88)
Back
Example 4-1B
212
4.2 Deduction of electronic structure from
ionization enthalpies (p.94)
Check Point 4-2

• Given the successive ionization enthalpies of
  boron, plot a graph of the logarithm of
  successive ionization enthalpies of boron against
  the number of electrons removed. Comment on the
  graph obtained.
• Successive I.E. (in kJ mol-1) 800, 2400, 3700,
  25000, 32800

Answer
213
4.2 Deduction of electronic structure from
ionization enthalpies (p.94)
Check Point 4-2
214
4.2 Deduction of electronic structure from
ionization enthalpies (p.94)
Check Point 4-2
(b) Give a sketch of the logarithm of successive
ionization enthalpies of potassium against no. of
electrons removed. Explain your sketch.
Answer
215
4.2 Deduction of electronic structure from
ionization enthalpies (p.94)
Check Point 4-2
216
4.2 Deduction of electronic structure from
ionization enthalpies (p.94)
Check Point 4-2

• There is always a drastic increase in ionization
  enthalpy whenever electrons are removed from a
  completely filled electron shell. Explain
  briefly.

Answer
(c) A completely filled electron shell has extra
stability. Once an electron is removed, the
stable electronic configuration will be
destroyed. Therefore, a larger amount of energy
is required to remove an electron from such a
stable electronic configuration.
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217
4.3 The Wave-mechanical model of the atom (p.97)
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Check Point 4-3

• What are the limitations of Bohrs atomic model?
• Explain the term dual nature of electrons.
• (c) For principal quantum number 4, how many
  sub-shells are present? What are their symbols?

Answer
218
4.4 Atomic orbitals (p.101)
Check Point 4-4

• Distinguish between the terms orbit and orbital.
• Sketch the pictorial representations of an s
  orbital and a p orbital. What shapes are they?

Answer
219
4.4 Atomic orbitals (p.101)
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Check Point 4-4

• How do the 1s and 2s orbitals differ from each
  other?
• (d) How do the 2p orbitals differ from each
  other?

(c) Both 1s and 2s orbitals are spherical in
shape, but the 2s orbital consists of a region of
zero probability of finding the electron known as
a nodal surface.
(d) There are three types of p orbitals. All are
dumb-bell in shape. They are aligned in three
different spatial orientations designated as x, y
and z. Hence, the 2p orbitals are designated as
2px, 2py and 2pz.
Answer