new way chemistry for hong kong a-level book 1 1 chapter 5 electronic configurations and the...
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New Way Chemistry for Hong Kong A-Level Book 1
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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)
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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
5.1 Relative Energies of Orbitals (SB p.112)
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ClassworkClasswork
Draw the electron-in-box diagrams and write the electronic configurations for the first 20 elements in the Periodic Table.
5.1 Relative Energies of Orbitals (SB p.114)
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1H1s
Element Electron-in-box Diagram Electronic Configuration
8O1s 2s 2p
1s 2s 2p 3s11Na 1s22s22p63s2
1s22s22p4
1s1
5.1 Relative Energies of Orbitals (SB p.114)
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19K
1s 2s 2p
3d
3s
4s
3p
3d 4s
[Ar]
Can be simplified as:
5.1 Relative Energies of Orbitals (SB p.114)
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ClassworkClasswork
Draw the electron-in-box diagrams and write the electronic configurations for the elements with atomic numbers from 21 to 30.
5.1 Relative Energies of Orbitals (SB p.114)
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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
5.1 Relative Energies of Orbitals (SB p.114)
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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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5.2 Relative Electronic Configurations of Elements (p. 115)
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
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Represented by ‘Electron-in-boxes’ Diagrams5.2 Relative Electronic Configurations of Elements (p. 117)
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5.2 Relative Electronic Configurations of Elements (p. 117)
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d-block
p-block
f-block
s-block
5.3 The Periodic Table (p. 118)
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Ionization Enthalpies of Elements5.4 Ionization Enthalpies of Elements (p. 120)
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Ionization Enthalpy across a Period5.4 Ionization Enthalpies of Elements (p. 122)
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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.
5.4 Ionization Enthalpies of Elements (p. 122)
•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.
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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.
5.4 Ionization Enthalpies of Elements (p. 123)
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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.
5.4 Ionization Enthalpies of Elements (p. 123)
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Q: Explain why there is large drop of I.E. between periods.Q: Explain why there is large drop of I.E. between periods.
5.4 Ionization Enthalpies of Elements (p. 123)
• 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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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.
• Thus the atomic radius increases, making the nucleus less effective in holding the s electron in the outer shell
• 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
5.4 Ionization Enthalpies of Elements (p. 123)
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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.
5.4 Ionization Enthalpies of Elements (p. 123)
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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.
5.4 Ionization Enthalpies of Elements (p. 123)
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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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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
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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5.5 Variation of Successive Ionization Enthalpies with Atomic Numbers (p. 126)
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Atomic size of elements5.6 Atomic Size of Elements (p. 128)
…..
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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.
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+11
Sodium atom Na(2,8,1)
5.6 Atomic Size of Elements (p. 128)
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+9
Sodium atom Na(2,8,1)
5.6 Atomic Size of Elements (p. 128)
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+1
Sodium atom Na(2,8,1)
Effective nuclear charge = +1
5.6 Atomic Size of Elements (p. 128)
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+12
Magnesium Mg(2,8,2)
5.6 Atomic Size of Elements (p. 128)
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+10
Magnesium Mg(2,8,2)
5.6 Atomic Size of Elements (p. 128)
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+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.
5.6 Atomic Size of Elements (p. 128)
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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.
5.6 Atomic Size of Elements (p. 129)
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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:
5.6 Atomic Size of Elements (p. 129)