Electronic Configurations

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
Relative energies of orbitals
5.1 Relative energies of orbitals (SB p.106)
4
Building up of electronic configurations
5.1 Relative energies of orbitals (SB p.106)
5
5.1 Relative energies of orbitals (SB p.106)
Aufbau principle states that electrons will enter
the possible orbitals in the order of ascending
energy.
Paulis exclusion principle states that electrons
occupying the same orbital must have opposite
spins.
Hunds rule (Rule of maximum multiplicity) states
that electrons must occupy each energy level
singly before pairing takes place (because of
their mutual repulsion).
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Electronic Configurations of Elements
7
Represented by notations
5.2 Electronic configurations of elements (SB
p.108)
8
5.2 Electronic configurations of elements (SB
p.109)
Represented by notations
9
5.2 Electronic configurations of elements (SB
p.109)
Represented by notations
10
Represented by electrons-in-boxes diagrams
5.2 Electronic configurations of elements (SB
p.110)
11
5.2 Electronic configurations of elements (SB
p.110)
12
The Periodic Table
13
The Periodic Table
5.3 The Periodic Table (SB p.112)
14
5.3 The Periodic Table (SB p.112)
s-block
p-block
d-block
f-block
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5.3 The Periodic Table (SB p.112)
Let's Think 1
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Ionization Enthalpies of Elements
17
Ionization enthalpies of elements
5.4 Ionization enthalpies of elements (SB p.115)
The first ionization enthalpies of the first 36
elements
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5.4 Ionization enthalpies of elements (SB p.116)
The first ionization enthalpies generally
decrease down a group and increases across a
period
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Ionization enthalpy across a period
5.4 Ionization enthalpies of elements (SB p.116)
20
5.4 Ionization enthalpies of elements (SB p.116)
Q Explain why there is a general increase in the
ionization energy across a period.

• Moving across a period, there is an increase in
  the nuclear attraction due to the addition of
  proton in the nucleus.
• The added electron is placed in the same quantum
  shell. It is only poorly shielded 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.
• The increase in the effective nuclear charge
  causes a decrease in the atomic radius.

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5.4 Ionization enthalpies of elements (SB p.117)
22
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is a trough at Boron(B) in
Period 2.

• e.c. of Be 1s22s2e.c. of B 1s22s22p1
• It is easier to remove the less penetrating
  p-electron from B than to remove a s electron
  from a stable fully-filled 2s subshell in Be.

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5.4 Ionization enthalpies of elements (SB p.117)
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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 1s22s22p3e.c. of O
  1s22s22p4
• It is more difficult to remove an electron from
  the halfly-filled 2p subshell of P, which has
  extra stability.
• After the removal of a p electron, a stable
  half-filled 2 p subshell can be obtained for Q.

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5.4 Ionization enthalpies of elements (SB p.117)
26
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is large drop of I.E.
between periods.

• The element at the end of a period has a stable
  octet structure. Much energy is required to
  remove an electron from it as this will disturb
  the stable structure.
• The element at the beginning of the next period
  has one extra s electron in an outer quantum
  shell. Although there is also an increase in the
  nuclear charge, it is offset very effectively by
  the screening effect of the inner shell
  electrons.
• Thus the atomic radius increases, making the
  nucleus less effective in holding the s electron
  in the outer shell

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5.4 Ionization enthalpies of elements (SB p.117)
28
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why there is drop of I.E. down a group.

• In 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.
• Thus the atomic radius increases, making the
  nucleus less effective in holding the s electron
  in the outer shell

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5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why successive ionization energies
increase.

• It is more difficult to remove electron(negatively
  charged) from higher positively charged ions.

30
5.4 Ionization enthalpies of elements (SB p.117)
Q Explain why successive ionization energy curve
follows the same pattern as the last one, but is
shifted by one unit of atomic number to the right.

• It is because the electronic configuration of AZ
  is the same as Az-1.

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Variation of Successive Ionization Enthalpies
with Atomic Numbers
32
Successive Ionization Enthalpies of the first 20
elements
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 119)
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5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 119)
34
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 120)
Variation of the first, second and third
ionization enthalpies of the first 20 elements
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Atomic Size of Elements
36
Atomic size of elements
5.6 Atomic size of elements (p. 122)
37
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.
• The added electron is placed in the same quantum
  shell. It is only poorly 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.

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5.6 Atomic size of elements (p. 122)
11
Sodium atom Na (2,8,1)
39
5.6 Atomic size of elements (p. 122)
9
Sodium atom Na (2,8,1)
40
5.6 Atomic size of elements (p. 122)
1
Effective nuclear charge 1
Sodium atom Na (2,8,1)
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5.6 Atomic size of elements (p. 122)
12
Magnesium atom Mg (2,8,2)
42
5.6 Atomic size of elements (p. 122)
10
Magnesium atom Mg (2,8,2)
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5.6 Atomic size of elements (p. 122)
By similar argument, effective nuclear charge
2 for a Mg atom.
2
Magnesium atom Mg (2,8,2)
Thus effective nuclear charge increases across a
period.
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5.6 Atomic size of elements (p. 122)
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5.6 Atomic size of elements (p. 122)
Q Explain why the atomic radius increases down a
group.

• 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 one more
  electron shell occupied which lies at a greater
  distance from the nucleus.

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5.6 Atomic size of elements (p. 122)
Remarks
Effective nuclear charge can only be applied to
make comparison between atoms in the same period.
Never apply effective nuclear charge to atoms in
the same group.
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The END
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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
49
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
50
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.
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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
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5.3 The Periodic Table (SB p.114)
Back
Check Point 5-3
53
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.

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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.

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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.
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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
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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.

Answer
58
5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 121)
Example 5-5
59
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.
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5.5 Variation of successive ionization enthalpies
with atomic numbers (p. 122)
Check Point 5-5

• Give the electrons-in-boxes diagram of 26Fe.
• 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
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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
62
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.

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

• (i)

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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
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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.
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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.