1. Electronic Configurations
Chapter 5
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
5.1 Relative Energies of Orbitals (SB p.112)
Relative Energies of Orbitals

3. Building up of Electronic Configurations
5.1 Relative Energies of Orbitals (SB p.112)
Building up of Electronic Configurations

4. 5.1 Relative Energies of Orbitals (SB p.112)
Aufbau principle states that electrons will enter the possible orbitals in the order of ascending energy.
Pauli’s exclusion principle states that no two electrons in the same atom can have identical values for all four sets of quantum numbers.
Hund’s rule (Rule of maximum multiplicity) states that electrons must occupy each energy level singly before pairing takes place (because of their mutual repulsion) and only then does pairing occur.
Carbon 1s 2s 2p

5. 5.1 Relative Energies of Orbitals (SB p.114)
Classwork
Draw the electron-in-box diagrams and write the electronic configurations for the first 20 elements in the Periodic Table.

6. 1H 8O 11Na 1s1 1s22s22p4 1s22s22p63s2 Element Electron-in-box Diagram
5.1 Relative Energies of Orbitals (SB p.114)
Element
Electron-in-box Diagram
Electronic Configuration 1H
1s1 1s 8O 1s 2s 2p
1s22s22p4 1s 2s 2p 3s
11Na
1s22s22p63s2

7. [Ar] 19K 1s 2s 2p 3d 3s 4s 3p Can be simplified as: 3d 4s
5.1 Relative Energies of Orbitals (SB p.114)
19K 1s 2s 2p 3d 3s 4s 3p
Can be simplified as: 3d 4s
[Ar]

8. 5.1 Relative Energies of Orbitals (SB p.114)
Classwork
Draw the electron-in-box diagrams and write the electronic configurations for the elements with atomic numbers from 21 to 30.

9. [Ar] 21Sc [Ar] 24Cr [Ar] 29Cu 3d 4s 3d 4s
5.1 Relative Energies of Orbitals (SB p.114) 3d 4s
[Ar]
21Sc 3d 4s
[Ar]
24Cr
Halfly-filled subshell  extra stability 3d 4s
[Ar]
29Cu
Fully-filled subshell  extra stability

10. Electronic Configurations of Isolated Atoms
5.2 Relative Electronic Configurations of Elements (p. 114)
Electronic Configurations of Isolated Atoms
Atomic no.
Element
Symbol
Arrangement of electrons in shells
Electronic configuration
“Standard form”
“Abbreviated form” 1 2 3 4 5 6 7 8
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen H He Li Be B C N O
2,1
2,2
2,3
2,4
2,5
2,6
1s1
1s2
1s22s1
1s22s2
1s22s22p1
1s22s22p2
1s22s22p3
1s22s22p4
[He]2s1
[He]2s2
[He]2s22p1
[He]2s22p2
[He]2s22p3
[He]2s22p4

11. Electronic Configurations of Isolated Atoms
5.2 Relative Electronic Configurations of Elements (p. 115)
Electronic Configurations of Isolated Atoms
Atomic no.
Element
Symbol
Arrangement of electrons in shells
Electronic configuration
“Standard form”
“Abbreviated form” 9 10 11 12 13 14 15 16
Fluorine
Neon
Sodium
Magnesium
Aluminium
Silicon
Phoshporus
Sulphur F Ne Na Mg Al Si P S
2,7
2,8
2,8,1
2,8,2
2,8,3
2,8,4
2,8,5
2,8,6
1s22s22p5
1s22s22p6
1s22s22p63s1
1s22s22p63s2
1s22s22p63s23p1
1s22s22p63s23p2
1s22s22p63s23p3
1s22s22p63s23p4
[He]2s22p5
[He]2s22p6
[Ne]3s1
[Ne]3s2
[Ne]3s23p1
[Ne]3s23p2
[Ne]3s23p3
[Ne]3s23p4

12. Electronic Configurations of Isolated Atoms
5.2 Relative Electronic Configurations of Elements (p. 115)
Electronic Configurations of Isolated Atoms
Atomic no.
Element
Symbol
Arrange-ment of electrons in shells
Electronic configuration
“Standard form”
“Abbreviat-ed form” 17 18 19 20
Chlorine
Argon
Potassium
Calcium Cl Ar K Ca
2,8,7
2,8,8
2,8,8,1
2,8,8,2
1s22s22p63s23p5
1s22s22p63s23p6
1s22s22p63s23p64s1
1s22s22p63s23p64s2
[Ne]3s23p5
[Ne]3s23p6
[Ar]4s1
Ar]4s2

13. Represented by ‘Electron-in-boxes’ Diagrams
5.2 Relative Electronic Configurations of Elements (p. 117)
Represented by ‘Electron-in-boxes’ Diagrams

14. 5.2 Relative Electronic Configurations of Elements (p. 117)

15. 5.3 The Periodic Table (p. 118)
The Periodic Table

16. 5.3 The Periodic Table (p. 118)
p-block
s-block
d-block
f-block

17. 5.3 The Periodic Table (p. 119)

18. Ionization Enthalpies of Elements
5.4 Ionization Enthalpies of Elements (p. 120)
Ionization Enthalpies of Elements

19. 5.4 Ionization Enthalpies of Elements (p. 121)

20. Ionization Enthalpy across a Period
5.4 Ionization Enthalpies of Elements (p. 122)
Ionization Enthalpy across a Period

21. 5.4 Ionization Enthalpies of Elements (p. 122)
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.

22. 5.4 Ionization Enthalpies of Elements (p. 123)

23. Q: Explain why there is a trough at Boron(B) in Period 2.
5.4 Ionization Enthalpies of Elements (p. 123)
Q: Explain why there is a trough at Boron(B) in Period 2.
e.c. of Be : 1s22s2 e.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.

24. 5.4 Ionization Enthalpies of Elements (p. 123)

25. Q: Explain why there is a trough at Oxygen(O) in Period 2.
5.4 Ionization Enthalpies of Elements (p. 123)
Q: Explain why there is a trough at Oxygen(O) in Period 2.
e.c. of N : 1s22s22p3 e.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.

26. 5.4 Ionization Enthalpies of Elements (p. 123)

27. Q: Explain why there is large drop of I.E. between periods.
5.4 Ionization Enthalpies of Elements (p. 123)
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

28. 5.4 Ionization Enthalpies of Elements (p. 123)

29. Q: Explain why there is drop of I.E. down a group.
5.4 Ionization Enthalpies of Elements (p. 123)
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

30. Q: Explain why successive ionization energies increase.
5.4 Ionization Enthalpies of Elements (p. 123)
Q: Explain why successive ionization energies increase.
It is more difficult to remove electron(negatively charged) from higher positively charged ions.

31. It is because the electronic configuration of AZ+ is the same as Az-1.
5.4 Ionization Enthalpies of Elements (p. 123)
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.

32. Successive Ionization Enthalpies with Atomic Number
5.5 Variation of Successive Ionization Enthalpies with Atomic Numbers (p. 124)
Successive Ionization Enthalpies with Atomic Number
Atomic
number Element
ΔH I.E. (kJ mol-1)
1 st
2nd
3rd
4th 1 2 3 4 5 6 7 8 9 10 H He Li Be B C N O F Ne
1 310
2 370
519
900
799
1 090
1 400
1 680
2 080
5 250
7 300
1 760
2 420
2 350
2 860
3 390
3 370
3 950
11 800
14 800
3 660
4 610
4 509
5 320
6 040
6 150
21 000
25 000
6 220
7 480
7 450
8 410
9 290

33. 5.5 Variation of Successive Ionization Enthalpies with Atomic Numbers (p. 124)
number Element
ΔH I.E. (kJ mol-1)
1 st
2nd
3rd
4th 11 12 13 14 15 16 17 18 19 20 Na Mg Al SI P S Cl Ar K Ca
494
736
577
786
1 060
1 000
1 260
1 520
418
590
4 560
1 450
1 820
1 580
1 900
2 260
2 300
2 660
3 070
1 150
6 940
7 740
2 740
3 230
2 920
3 390
3 850
3 950
4 600
4 940
9 540
10 500
11 600
4 360
4 960
4 540
5 150
5 77
5 860
6 480

34. 5.5 Variation of Successive Ionization Enthalpies with Atomic Numbers (p. 126)

35. Atomic size of elements
5.6 Atomic Size of Elements (p. 128)
Atomic size of elements
…..

36. Q: Explain why the atomic radius decreases across a period.
5.6 Atomic Size of Elements (p. 128)
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.

37. 5.6 Atomic Size of Elements (p. 128)
+11
Sodium atom Na
(2,8,1)

38. 5.6 Atomic Size of Elements (p. 128)+9
Sodium atom Na
(2,8,1)

39. Effective nuclear charge = +1
5.6 Atomic Size of Elements (p. 128) +1
Effective nuclear charge = +1
Sodium atom Na
(2,8,1)

40. 5.6 Atomic Size of Elements (p. 128)
+12
Magnesium Mg
(2,8,2)

41. 5.6 Atomic Size of Elements (p. 128)
+10
Magnesium Mg
(2,8,2)

42. By similar argument, effective nuclear charge = +2 for a Mg atom.
5.6 Atomic Size of Elements (p. 128) +2
Magnesium Mg
(2,8,2)
By similar argument, effective nuclear charge = +2 for a Mg atom.
Thus effective nuclear charge increases across a period.

43. 5.6 Atomic Size of Elements (p. 129)

44. Q: Explain why the atomic radius increases down a group.
5.6 Atomic Size of Elements (p. 129)
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.

45. Never apply effective nuclear charge to atoms in the same group.
5.6 Atomic Size of Elements (p. 129)
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.

46. The END