Electronic Configurations and the Periodic Table

1. 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. 5.1 Relative Energies of Orbitals

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

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

5. 5.1 Relative energies of orbitals (SB p.106) Carbon Check Point 5-1 1s 2s 2p 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).

6. 5.2 Electronic Configurations of Elements

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

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

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

12. 5.3 The Periodic Table

13. 5.3 The Periodic Table (SB p.112) The Periodic Table

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) Check Point 5-3 Let's Think 1

16. 5.4 Ionization Enthalpies of Elements

17. 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. 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. 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 theless penetrating p-electron from Bthan to remove as electron from a stable fully-filled 2ssubshell in Be.

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

24. 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 thehalfly-filled 2p subshell of P,which has extra stability. • After the removal of a p electron, a stable half-filled 2 psubshell can be obtained for Q.

25. 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 octetstructure. 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 veryeffectively by the screening effect of the inner shellelectrons. • 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. 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 veryeffectively by the screening effect of the inner shellelectrons. • Thus the atomic radius increases, making the nucleus less effective in holding the s electron in the outer shell

29. 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) Check Point 5-4 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 thesame as Az-1.

31. 5.5 Variation of Successive Ionization Enthalpies with Atomic Numbers

32. 5.5 Variation of successive ionization enthalpies with atomic numbers (p. 119) ΔH I.E. (kJ mol-1) Element Atomic number 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 1 310 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 Successive Ionization Enthalpies of the first 20 elements

33. 5.5 Variation of successive ionization enthalpies with atomic numbers (p. 119) Atomic number ΔH I.E. (kJ mol-1) Element 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. 120) Example 5-5 Check Point 5-5 Variation of the first, second and third ionization enthalpies of the first 20 elements

35. 5.6 Atomic Size of Elements

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

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.

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. 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. 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. 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 screeningeffect of the inner shell electrons. • Moving down a group, an atom would have one moreelectron shell occupied which lies at a greater distance from the nucleus.

46. 5.6 Atomic size of elements (p. 122) Check Point 5-6 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.

47. The END

48. 5.1 Relative energies of orbitals (SB p.108) • Nitrogen: 1s22s22p3 • (b) Sodium: 1s22s22p63s1 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) • Silicon: [Ne]3s23p3 • (b) Copper: [Ar]3d104s1 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.