New Way Chemistry for Hong Kong A-Level Book 1 1 Electronic Configurations and the Periodic Table...

New Way Chemistry for Hong Kong A-Level Book 1

1

Electronic Configurations Electronic Configurations

and the Periodic Tableand the Periodic Table

5.15.1 Relative Energies of Orbitals Relative Energies of Orbitals

5.25.2 Electronic Configurations of Elements Electronic Configurations of Elements

5.35.3 The Periodic TableThe Periodic Table

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

5.45.4 Atomic Size of ElementsAtomic Size of Elements

55


2

5.5.11Relative EnergieRelative Energie

s of Orbitalss of Orbitals


3

Relative energies of orbitalsRelative energies of orbitals5.1 Relative energies of orbitals (SB p.106)


4

Building up of electronic Building up of electronic configurationsconfigurations

5.1 Relative energies of orbitals (SB p.106)


5

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 electrons occupying the same orbital must have opposite spins.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).

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 5-1Check Point 5-1



6

5.5.22 Electronic Electronic ConfigurationConfigurations of Elementss of Elements


7

Represented by notationsRepresented by notations

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

5.2 Electronic configurations of elements (SB p.108)


8

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




9

Atomic no.

Element Symbol Arrange-ment

of electrons in s

hells

Electronic configuration“Standard form” “Abbreviat-

ed 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




10

Represented by ‘electrons-in-boxes’ Represented by ‘electrons-in-boxes’ diagramsdiagrams



11

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


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5.5.33 The Periodic The Periodic

TableTable


13

The Periodic TableThe Periodic Table

5.3 The Periodic Table (SB p.112)


14

d-block

p-block

f-block

s-block



15


Check Point 5-3Check Point 5-3


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5.5.44 Ionization Ionization Enthalpies of Enthalpies of

ElementsElements


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Ionization enthalpies of elementsIonization enthalpies of elements5.4 Ionization enthalpies of elements (SB p.115)

The first ionization enthalpies

of the first 36

elements


18

The first ionization enthalpies generally decrease down a

group and increases across a period

5.4 Ionization enthalpies of elements (SB p.116)


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Ionization enthalpy across a Ionization enthalpy across a periodperiod



20

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.



21



22

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.



23



24

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.



25



26

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



27



28

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.




29

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 A

Z+ is the same as Az-1.

• It is because the electronic configuration of A

Z+ 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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5.5.55 Variation of Variation of

Successive Successive Ionization Ionization

Enthalpies with Enthalpies with Atomic NumbersAtomic Numbers


32

Successive Ionization Enthalpies of Successive Ionization Enthalpies of the first 20 elementsthe first 20 elements

5.5 Variation of successive ionization enthalpies with atomic numbers (p. 119)

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

1 310

1 680

2 080

H

He

Li

Be

B

C

N

O

F

Ne

1

2

3

4

5

6

7

8

9

10

4th3rd2nd1 st

ΔH I.E. (kJ mol-1)Atomic

number

Element


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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

4th3rd2nd1 st

ΔH I.E. (kJ mol-1)Atomic

number

Element



34


Variation of the first, second

and third ionization

enthalpies of the first 20 elements


Example 5-5Example 5-5


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5.5.66 Atomic Size Atomic Size

of Elementsof Elements


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Atomic size of elementsAtomic size of elements

…..

5.6 Atomic size of elements (p. 122)


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



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+11

Sodium atom Na(2,8,1)



39

+9




40

+1

Effective nuclear charge = +1




41

+12

Magnesium atom Mg(2,8,2)



42

+10




43

+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.




44



45

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


48

Write the electronic configurations and draw “electrons-in –boxes” diagrams for

(a) nitrogen; and

(b) sodium.

Back

Answer


(a) Nitrogen: 1s22s22p3

(b) Sodium: 1s22s22p63s1


49

Give the electronic configuration by notations and “electrons-in-boxes” diagrams in the abbreviated form for the following elements.

(a) silicon; and

(b) copper.

Back

Answer


(a) Silicon: [Ne]3s23p3

(b) Copper: [Ar]3d104s1


50

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?

Back

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.

Answer



51


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


Back


53

(a)Give four main factors that affect the magnitude of ionization enthalpy of an atom.

Answer


(a) 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

(b)Explain why Group 0 elements have extra high first ionization enthalpies and their decreasing trend down the group.

Answer


(b) 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.


55

(c)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.

Back


56


For the element 126C,

(a)(i) write its electronic configuration by notation.

(ii) write its electronic configuration by “electrons-in- boxes” diagram. Answer(a) (i) 1s22s22p2

(ii)


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(b)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. Answer

1st 2nd 3rd 4th 5th 6thI.E. (kJ

mol-1)

1090 2350 4610 6220 37800 47000


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(b) (i)


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(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.

Back


60


(a) Give the “electrons-in-boxes” diagram of 26Fe.

(b)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+.

Answer

(a) Fe :

(b) Fe2+ :

Fe3+ :


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(c) Which ion is more stable, Fe2+ or Fe3+? Explain briefly.

Answer

(c) Fe3+ ion is more stable because the 3d orbital is exactly half-filled which gives the electronic configuration extra stability.



62

(d)Given the successive ionization enthalpies of Fe:

(i) plot a graph of successive ionization enthalpies in logarithm scale against the number of electron

s removed;

(ii) state the difference of the plot from that of carbon as shown in P. 121.

Answer


1st 2nd 3rd 4th 5th 6thI.E. (kJ

mol-1)

762 1560 2960 5400 7620 10100


63


(d) (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


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(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 su

ccessive ionization enthalpies of iron.

Back


65

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 pro

gressively 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

Explain the following:

(b) The atomic radius increases down Group I metals.


Back

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.

New Way Chemistry for Hong Kong A-Level Book 1 1 Electronic Configurations and the Periodic Table...

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