1. Electronic Configurations5
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

3. Relative energies of orbitals
5.1 Relative energies of orbitals (SB p.106)
Relative energies of orbitals

4. Building up of electronic configurations
5.1 Relative energies of orbitals (SB p.106)
Building up of electronic configurations

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.
Pauli’s exclusion principle states that electrons occupying the same orbital must have opposite spins.
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).
Carbon 1s 2s 2p
Check Point 5-1

6. Electronic Configurations of Elements
5.2
Electronic Configurations of Elements

7. Represented by notations
5.2 Electronic configurations of elements (SB p.108)
Represented by notations
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

8. Arrangement of electrons in shells Electronic configuration
5.2 Electronic configurations of elements (SB p.109)
Represented by notations
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

9. Arrange-ment of electrons in shells Electronic configuration
5.2 Electronic configurations of elements (SB p.109)
Represented by notations
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

10. Represented by ‘electrons-in-boxes’ diagrams
5.2 Electronic configurations of elements (SB p.110)
Represented by ‘electrons-in-boxes’ diagrams

11. Check Point 5-2
5.2 Electronic configurations of elements (SB p.110)

12. 5.3
The Periodic Table

13. 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

15. 5.3 The Periodic Table (SB p.112)
Let's Think 1
Check Point 5-3

16. Ionization Enthalpies of Elements
5.4
Ionization Enthalpies of Elements

17. Ionization enthalpies of elements
5.4 Ionization enthalpies of elements (SB p.115)
Ionization enthalpies of elements
The first ionization enthalpies of the first 36 elements

18. 5.4 Ionization enthalpies of elements (SB p.116)
The first ionization enthalpies generally decrease down a group and increases across a period

19. Ionization enthalpy across a period
5.4 Ionization enthalpies of elements (SB p.116)
Ionization enthalpy across a period

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.

21. 5.4 Ionization enthalpies of elements (SB p.117)

22. Q: Explain why there is a trough at Boron(B) in Period 2.
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 : 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.

23. 5.4 Ionization enthalpies of elements (SB p.117)

24. Q: Explain why there is a trough at Oxygen(O) in Period 2.
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 : 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.

25. 5.4 Ionization enthalpies of elements (SB p.117)

26. Q: Explain why there is large drop of I.E. between periods.
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

27. 5.4 Ionization enthalpies of elements (SB p.117)

28. Q: Explain why there is drop of I.E. down a group.
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

29. Q: Explain why successive ionization energies increase.
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.
Check Point 5-4

31. Variation of Successive Ionization Enthalpies with Atomic Numbers
5.5
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)
Successive Ionization Enthalpies of the first 20 elements
21 000
25 000
6 220
7 480
7 450
8 410
9 290
11 800
14 800
3 660
4 610
4 509
5 320
6 040
6 150
5 250
7 300
1 760
2 420
2 350
2 860
3 390
3 370
3 950
1 310
2 370
519
900
799
1 090
1 400
1 680
2 080 H He Li Be B C N O F Ne 1 2 3 4 5 6 7 8 9 10
4th
3rd
2nd
1 st
ΔH I.E. (kJ mol-1)
Atomic
number
Element

33. 5.5 Variation of successive ionization enthalpies with atomic numbers (p. 119)
9 540
10 500
11 600
4 360
4 960
4 540
5 150
5 77
5 860
6 480
6 940
7 740
2 740
3 230
2 920
3 390
3 850
3 950
4 600
4 940
4 560
1 450
1 820
1 580
1 900
2 260
2 300
2 660
3 070
1 150
494
736
577
786
1 060
1 000
1 260
1 520
418
590 Na Mg Al SI P S Cl Ar K Ca 11 12 13 14 15 16 17 18 19 20
4th
3rd
2nd
1 st
ΔH I.E. (kJ mol-1)
Atomic
number
Element

34. 5.5 Variation of successive ionization enthalpies with atomic numbers (p. 120)
Example 5-5
Check Point 5-5
Variation of the first, second and third ionization enthalpies of the first 20 elements

35. Atomic Size of Elements
5.6
Atomic Size of Elements

36. Atomic size of elements
5.6 Atomic size of elements (p. 122)
Atomic size of elements
…..

37. Q: Explain why the atomic radius decreases across a period.
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.

38. 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. Effective nuclear charge = +1
5.6 Atomic size of elements (p. 122) +1
Effective nuclear charge = +1
Sodium atom Na
(2,8,1)

41. 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)

43. By similar argument, effective nuclear charge = +2 for a Mg atom.
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.

44. 5.6 Atomic size of elements (p. 122)

45. Q: Explain why the atomic radius increases down a group.
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.

46. 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.
Check Point 5-6

47. The END

48. Check Point 5-1 Answer Back
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
Nitrogen: 1s22s22p3
(b) Sodium: 1s22s22p63s1

49. Check Point 5-2 Answer Back
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
Silicon: [Ne]3s23p3
(b) Copper: [Ar]3d104s1

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.

51. 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

52. 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.

54. 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.
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.
Answer

55. Check Point 5-4 Answer Back
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.

56. Example 5-5 Answer For the element 126C,
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
(i) 1s22s22p2
(ii)

57. 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.
1st
2nd
3rd
4th
5th
6th
I.E. (kJ mol-1)
1090
2350
4610
6220
37800
47000
Answer

58. 5.5 Variation of successive ionization enthalpies with atomic numbers (p. 121)
Example 5-5
(i)

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.

60. Check Point 5-5 Answer Give the “electrons-in-boxes” diagram of 26Fe.
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+.
Fe :
Fe2+ :
Fe3+ :
Answer

61. 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.
1st
2nd
3rd
4th
5th
6th
I.E. (kJ mol-1)
762
1560
2960
5400
7620
10100
Answer

63. Number of electrons removed
5.5 Variation of successive ionization enthalpies with atomic numbers (p. 122)
Check Point 5-5
(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

64. 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

65. Check Point 5-6 Answer Explain the following:
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.

66. Check Point 5-6 Answer Back Explain the following:
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.