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

Relative Energies of Orbitals 5.1 Relative Energies of Orbitals (SB p.112) Relative Energies of Orbitals

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

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

5.1 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.

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

[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]

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.

[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

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

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

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

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

5.2 Relative Electronic Configurations of Elements (p. 117)

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

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

5.3 The Periodic Table (p. 119)

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

5.4 Ionization Enthalpies of Elements (p. 121)

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

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.

5.4 Ionization Enthalpies of Elements (p. 123)

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.

5.4 Ionization Enthalpies of Elements (p. 123)

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.

5.4 Ionization Enthalpies of Elements (p. 123)

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

5.4 Ionization Enthalpies of Elements (p. 123)

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

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.

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.

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

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

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

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

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.

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

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

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

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

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

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.

5.6 Atomic Size of Elements (p. 129)

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.

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.

The END