New Way Chemistry for Hong Kong A-Level Book 1 1 Chapter 5 Electronic Configurations and the Periodic Table 5.1 Relative Energies of Orbitals 5.2 Electronic.

New Way Chemistry for Hong Kong A-Level Book 1

1

Chapter 5Chapter 5Electronic Configurations Electronic Configurations

and the Periodic Tableand the Periodic Table5.1 5.1 Relative Energies of Orbitals Relative Energies of Orbitals

5.2 5.2 Electronic Configurations of Elements Electronic Configurations of Elements

5.3 5.3 The Periodic TableThe Periodic Table

5.45.4 Ionization Enthalpies of Elements Ionization Enthalpies of Elements5.55.5 Variation of Successive Ionization Ethalpies Variation of Successive Ionization Ethalpies with Atomic Numbers with Atomic Numbers

5.45.4 Atomic Size of Elements Atomic Size of Elements


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Relative Energies of Orbitals

5.1 Relative Energies of Orbitals (SB p.112)


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Building up of Electronic Configurations5.1 Relative Energies of Orbitals (SB p.112)


4

Aufbau principle states that electrons will enter the possible orbitals in the order of ascending energy.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.

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.

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

ClassworkClasswork

Draw the electron-in-box diagrams and write the electronic configurations for the first 20 elements in the Periodic Table.



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1H1s

Element Electron-in-box Diagram Electronic Configuration

8O1s 2s 2p

1s 2s 2p 3s11Na 1s22s22p63s2

1s22s22p4

1s1



7

19K

1s 2s 2p

3d

3s

4s

3p

3d 4s

[Ar]

Can be simplified as:



8

ClassworkClasswork

Draw the electron-in-box diagrams and write the electronic configurations for the elements with atomic numbers from 21 to 30.



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3d 4s

[Ar]21Sc

24Cr3d 4s

[Ar]

3d 4s

[Ar]29Cu

Halfly-filled subshell extra stability

Fully-filled subshell extra stability



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Electronic Configurations of Isolated Atoms5.2 Relative Electronic Configurations of Elements (p. 114)

Atomic no.

Element Symbol Arrangement of electrons in

shells

Electronic configuration

“Standard form”

“Abbreviated form”

12345678

HydrogenHeliumLithiumBerylliumBoronCarbonNitrogenOxygen

HHeLiBeBCNO

122,12,22,32,42,52,6

1s1

1s2

1s22s1

1s22s2

1s22s22p1

1s22s22p2

1s22s22p3

1s22s22p4

1s1

1s2

[He]2s1

[He]2s2

[He]2s22p1

[He]2s22p2

[He]2s22p3

[He]2s22p4


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5.2 Relative Electronic Configurations of Elements (p. 115)

Atomic no.

Element Symbol Arrangement of electrons

in shells

Electronic configuration

“Standard form” “Abbreviated form”

910111213141516

FluorineNeonSodiumMagnesiumAluminiumSiliconPhoshporusSulphur

FNeNaMgAlSiPS

2,72,82,8,12,8,22,8,32,8,42,8,52,8,6

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


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Atomic no.

Element Symbol Arrange-ment of electrons in

shells

Electronic configuration“Standard form” “Abbreviat-e

d form”

17181920

ChlorineArgonPotassiumCalcium

ClArKCa

2,8,72,8,82,8,8,12,8,8,2

1s22s22p63s23p5

1s22s22p63s23p6

1s22s22p63s23p64s1

1s22s22p63s23p64s2

[Ne]3s23p5

[Ne]3s23p6

[Ar]4s1

Ar]4s2

Electronic Configurations of Isolated Atoms


13

Represented by ‘Electron-in-boxes’ Diagrams5.2 Relative Electronic Configurations of Elements (p. 117)


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The Periodic Table5.3 The Periodic Table (p. 118)


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d-block

p-block

f-block

s-block

5.3 The Periodic Table (p. 118)


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5.3 The Periodic Table (p. 119)


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Ionization Enthalpies of Elements5.4 Ionization Enthalpies of Elements (p. 120)


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5.4 Ionization Enthalpies of Elements (p. 121)


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Ionization Enthalpy across a Period5.4 Ionization Enthalpies of Elements (p. 122)


21

Q: Explain why there is a general increase in the ionization energy across a period.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



23

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

• 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



25

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

• 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



27

Q: Explain why there is large drop of I.E. between periods.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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29

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


• 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.




30

Q: Explain why successive ionization energies increase.Q: Explain why successive ionization energies increase.

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

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



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• It is because the electronic configuration of AZ+ i

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

+ is the same as Az-1.

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.

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.



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Successive Ionization Enthalpies with Atomic Number

5.5 Variation of Successive Ionization Enthalpies with Atomic Numbers (p. 124)

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

1 310

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


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Atomic

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


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35

Atomic size of elements5.6 Atomic Size of Elements (p. 128)

…..


36

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

5.6 Atomic Size of Elements (p. 128)

• 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

+11

Sodium atom Na(2,8,1)



38

+9




39

+1


Effective nuclear charge = +1



40

+12

Magnesium Mg(2,8,2)



41

+10

Magnesium Mg(2,8,2)



42

+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.Thus effective nuclear charge increases across a period.



43



44

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

• 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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Effective nuclear charge can only be applied to make comparison between atoms in the same period.

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.Never apply effective nuclear charge to atoms in the same group.

Remarks:



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The END

New Way Chemistry for Hong Kong A-Level Book 1 1 Chapter 5 Electronic Configurations and the Periodic Table 5.1 Relative Energies of Orbitals 5.2 Electronic.

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