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ATOMIC ELECTRON CONFIGURATIONS

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CHAPTER 8 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY Sample Question Give the spectroscopic notation for P. 1s22s22p63s23p3 Determine the set of 4 quantum ... – PowerPoint PPT presentation

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Title: ATOMIC ELECTRON CONFIGURATIONS


1
CHAPTER 8
  • ATOMIC ELECTRON CONFIGURATIONS
  • AND
  • PERIODICITY

2
Chapter 8 Outline
  • Magnetic Spin
  • Periodic Trends
  • Effective nuclear charge
  • Atomic Radii Size
  • Ion Radii Size
  • Ionization Energy and Electron Affinities

3
ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY
4
Arrangement of Electrons in Atoms
  • Electrons in atoms are arranged as

SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
5
Arrangement of Electrons in Atoms
  • Each orbital can be assigned no more than 2
    electrons!

This is tied to the existence of a 4th quantum
number, the electron spin quantum number, ms.
6
Electron Spin Quantum Number, ms
Can be proven experimentally that an electron has
a spin. The two spin directions are given by ms
where ms 1/2 and -1/2.
7
Electron Spin Quantum Number
Diamagnetic substance NOT attracted to a
magnetic field. Paramagnetic substance is
attracted to a magnetic field. Substance has
unpaired electrons.
8
Paramagnetic -vs- Ferromagnetic
Unpaired electrons in an electric field alien
themselves to the applied field
Unpaired electrons in their normal, random
orientation. ALL metals behave this way, the
exception is Cobalt, the reason why Cobalt
magnets are so VERY expensive.
9
QUANTUM NUMBERS
  • n ---gt shell 1, 2, 3, 4, ...
  • l ---gt subshell 0, 1, 2, ... n - 1
  • ml ---gt orbital -l ... 0 ... l
  • ms ---gt electron spin 1/2 and -1/2

10
Pauli Exclusion Principle
  • No two electrons in the same atom can have the
    same set of 4 quantum numbers.
  • That is, each electron has a unique address set
    of 4 quantum numbers.

11
Electrons in Atoms
  • When n 1, then l 0
  • this shell has a single orbital (1s) to which
    2e- can be assigned.
  • When n 2, then l 0, 1
  • 2s orbital 2e-
  • three 2p orbitals 6e-
  • TOTAL 8e-

12
Electrons in Atoms
  • When n 3, then l 0, 1, 2
  • 3s orbital 2e-
  • three 3p orbitals 6e-
  • five 3d orbitals 10e-
  • TOTAL 18e-

13
Electrons in Atoms
  • When n 4, then l 0, 1, 2, 3
  • 4s orbital 2e-
  • three 4p orbitals 6e-
  • five 4d orbitals 10e-
  • seven 4f orbitals 14e-
  • TOTAL 32e-

14
Assigning Electrons to Atoms
  • Electrons generally assigned to orbitals of
    successively higher energy.
  • For H atoms, E - C(1/n2). E depends only on n.
  • For multi - electron atoms, energy depends on
    both n and l.

15
Figure 8.4
16
Assigning Electrons to Subshells
  • In the H atom all subshells of same n have same
    energy.
  • In many-electron atom
  • a) subshells increase in energy as value of n l
    increases.
  • b) for subshells of same n l, subshell with
    lower n is lower in energy.

17
Electron Filling OrderFigure 8.5
18
Writing Atomic Electron Configurations
  • Two ways of writing complete configurations.
  • One is called the spectroscopic notation.

19
Writing Atomic Electron Configurations
  • Two ways of writing complete configurations.
  • Other is called the orbital box notation.

One electron has n 1, l 0, ml 0, ms
1/2 Other electron has n 1, l 0, ml 0, ms
- 1/2
20
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21
Figure 8.7
22
Lithium
  • Group 1A
  • Atomic number 3
  • 1s22s1 ---gt 3 total electrons

23
Beryllium
  • Group 2A
  • Atomic number 4
  • 1s22s2 ---gt 4 total electrons

24
Boron
  • Group 3A
  • Atomic number 5
  • 1s2 2s2 2p1 ---gt 5 total electrons

25
Carbon
  • Group 4A
  • Atomic number 6
  • 1s2 2s2 2p2 ---gt 6 total electrons

Here we see HUNDS RULE. When placing
electrons in a set of orbitals having the same
energy, we place them singly as long as possible.
26
Nitrogen
  • Group 5A
  • Atomic number 7
  • 1s2 2s2 2p3 ---gt 7 total electrons

27
Oxygen
  • Group 6A
  • Atomic number 8
  • 1s2 2s2 2p4 ---gt 8 total electrons

28
Fluorine
  • Group 7A
  • Atomic number 9
  • 1s2 2s2 2p5 ---gt 9 total electrons

29
Neon
  • Group 8A
  • Atomic number 10
  • 1s2 2s2 2p6 ---gt 10 total electrons

Note that we have reached the end of the 2nd
period, and the 2nd shell is full!
30
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31
Electron Configurations
Noble gas configuration filled s and p
levels. Electron configs. are shortened by using
the noble gas configuration to represent the core
electrons. The remaining electrons are valence
electrons. Notations Examples
Boron Spectroscopic 1s22s22p1 Orbital Box
Noble Gas He 2s22p1 Valence 2s22p1
32
Quantum Numbers
If we use positive values of ml and ms first, the
set of four quantum numbers can be determined for
the last electron in for a given element.
This is called the last-electron-in
game!!!
33
Electron Configurations of p-Block Elements
34
Sodium
  • Group 1A
  • Atomic number 11
  • 1s2 2s2 2p6 3s1 or
  • neon core 3s1
  • Ne 3s1 (uses rare gas notation)
  • Note that we have begun a new period.
  • All Group 1A elements have corens1
    configurations.

35
Aluminum
  • Group 3A
  • Atomic number 13
  • 1s2 2s2 2p6 3s2 3p1
  • Ne 3s2 3p1

All Group 3A elements have core ns2 np1
configurations where n is the period number.
36
Phosphorus
  • Group 5A
  • Atomic number 15
  • 1s2 2s2 2p6 3s2 3p3
  • Ne 3s2 3p3

All Group 5A elements have core ns2 np3
configurations where n is the period number.
37
Calcium
  • Group 2A
  • Atomic number 20
  • 1s2 2s2 2p6 3s2 3p6 4s2
  • Ar 4s2
  • All Group 2A elements have corens2
    configurations where n is the period number.

38
Relationship of Electron Configuration and Region
of the Periodic Table
  • Gray s block
  • Orange p block
  • Green d block
  • Violet f block

39
Transition MetalsTable 8.4
  • All 4th period elements have the configuration
    argon nsx (n - 1)dy and so are d-block
    elements.

Chromium
Copper
Iron
40
Transition Element Configurations
3d orbitals used for Sc - Zn (Table 8.4)
41
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42
Figure 8.9
43
Lanthanides and Actinides
  • All these elements have the configuration
    core nsx (n - 1)dy (n - 2)fz and so are
    f-block elements.

44
Lanthanide Element Configurations
4f orbitals used for Ce - Lu and 5f for Th - Lr
(Table 8.2)
45
Ion Configurations
  • To form cations from elements remove 1 or more
    e- from subshell of highest n or highest (n
    l).
  • P Ne 3s2 3p3 ---gt P3 Ne 3s2 3p0 3e-

P3
46
Ion Configurations
  • For transition metals, remove ns electrons and
    then (n - 1) electrons.
  • Fe Ar 4s2 3d6 ---gt Fe2 Ar 4s0 3d6 2e-

Fe2
47
Ion Configurations
  • How do we know the configurations of ions?
  • Determine the magnetic properties of ions.
  • Ions with UNPAIRED ELECTRONS are PARAMAGNETIC.
  • Without unpaired electrons DIAMAGNETIC.

48
Sample Question
  • Give the electron configuration for Ni4.
  • Include the valence box diagram and state if
    the ion is paramagnetic or diamagnetic.

3d
Ni4

49
PERIODIC TRENDS
Movies on these later!!
50
General Periodic Trends
  • Atomic and ionic size
  • Ionization energy
  • Electron affinity

51
Atomic Size
  • Size goes UP when going down a group. See
    Figure 8.9.
  • Because electrons are added further from the
    nucleus, there is less attraction.
  • Size goes DOWN when going across a period.

52
Effective Nuclear Charge relating to Atom radii
  • The reason for the difference in energy for 2s
    and 2p subshells, for example, is the effective
    nuclear charge, Z.

53
Figure 8.6
Screening or Shielding effect helps explain
periodic trends in a period
54
Effective Nuclear Charge, Z
  • Z is the nuclear charge experienced by the
    outermost electrons.
  • Explains why E(2s) lt E(2p)
  • Z increases across a period owing to incomplete
    shielding by inner electrons.
  • Estimate Z by --gt Z - ( inner electrons)
  • Charge felt by 2s e- in Li Z 3 - 2 1
  • Be Z 4 - 2 2
  • B Z 5 - 2 3 and so on!

55
The nuclear charge increases
56
Atomic Size
  • Size decreases across a period owing to
    increase in Z. Each added electron feels a
    greater and greater charge.

57
Trends in Atomic Size
58
Sizes of Transition Elements
  • 3d subshell is inside the 4s subshell.
  • 4s electrons feel a more or less constant Z.
  • Sizes stay about the same and chemistries are
    similar!

59
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60
Ion Sizes
Does the size go up or down when losing an
electron to form a cation?
61
Ion Sizes
  • CATIONS are SMALLER than the atoms from which
    they come.
  • The electron/proton attraction has gone UP and so
    size DECREASES.

62
Ion Sizes
  • Does the size go up or down when gaining an
    electron to form an anion?

63
Ion Sizes
  • ANIONS are LARGER than the atoms from which they
    come.
  • The electron/proton attraction has gone DOWN and
    so size INCREASES.

64
Figure 8.13
65
Ion Sizes
  • For an isoelectronic series
  • Size decreases as the atomic number increases
  • Example N-3 to Al3

66
Redox Reactions
  • Why do metals lose electrons in their
    reactions?
  • Why does Mg form Mg2 ions and not Mg3?
  • Why do nonmetals take on electrons?

67
Ionization EnergySee Figure 8.12
  • IE energy required to remove an electron from
    an atom in the gas phase.
  • Mg (g) 738 kJ ---gt Mg (g) e-

A(g) ---gt A(g) e- I.E. ?E
68
Ionization Energy
  • IE energy required to remove an electron from
    an atom in the gas phase.
  • Mg (g) 738 kJ ---gt Mg (g) e-
  • Mg (g) 1451 kJ ---gt Mg2 (g) e-
  • Mg has 12 protons and only 11 electrons.
    Therefore, IE for Mg gt Mg.

69
Ionization Energy
  • Mg (g) 735 kJ ---gt Mg (g) e-
  • Mg (g) 1451 kJ ---gt Mg2 (g) e-
  • Mg2 (g) 7733 kJ ---gt Mg3 (g) e-
  • Energy cost is very high to dip into a shell
    of lower n. This is why ox. no. Group no.

70
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71
Trends in Ionization Energy
  • IE increases across a period because Z
    increases.
  • Metals lose electrons more easily than nonmetals.
  • Metals are good reducing agents.
  • Nonmetals lose electrons with difficulty.

72
Trends in Ionization Energy
  • IE decreases down a group
  • Because size increases.
  • Reducing ability generally increases down the
    periodic table.
  • See reactions of Li, Na, K

73
Figure 8.11
74
Periodic Trend in the Reactivity of Alkali Metals
with Water
Lithium
Sodium
Potassium
Here are those movies promised!
75
Electron Affinity
  • A few elements GAIN electrons to form anions.
  • Electron affinity is the energy involved when
    an atom Gains an electron.
  • A(g) e- ---gt A- (g) E.A. -DE

Note Both EAs and IEs are -.
Cl E.A. -349 kJ/Mol
76
Electron Affinity of Oxygen
EA - 141 kJ
  • DE is Exothermic because O has an affinity for
    an e-.

electron
77
Electron Affinity of Nitrogen
EA 0 kJ
  • DE is zero for N due to electron-electron
    repulsions.

electron
78
Trends in Electron Affinity
  • See Figure 8.12
  • Affinity for electron increases across a period
    (EA becomes more negative).
  • Affinity decreases down a group (EA becomes less
    negative).

79
Figure 8.12
80
Trends in Electron Affinity
  • See Appendix F
  • Affinity for electron increases across a period
    (EA becomes more negative).
  • Affinity decreases down a group (EA becomes less
    negative).

Atom EA F -328 kJ Cl -349 kJ Br -325 kJ I -295
kJ
These numbers are for the ion forming the atom.
81
Practice Problems
  • 1. Write the electron configurations
  • a. Orbital box notation for C
  • b. Spectroscopic notation for Mg
  • c. Noble Gas notation for Sc
  • d. Noble Gas notation for P
  • e. Valence notation for Pb

82
Practice Problems
2. Answer the following about the elements in
question 1 a. number of unpaired
electron's b. arrange according to size (small
to large) c. arrange according to ionization
energy (small to large) d. predict ions that
each elements will form e. determine the set of
4 quantum numbers for the last electron.
83
Practice Problems
3. Give the symbol of the element with the lowest
atomic number that has a. a half filled p
sublevel b. 3 filled 4d orbitals c. unpaired
electrons in 2 or more sublevels
84
Practice Problems Answers
1. a. b. 1s22s22p63s2 c. Ar 4s23d1 d. Ne
3s23p3 e. 6s26p2 or 6s24f145d106p2 2. a. 2, 0,
1, 3, 2 b. C, P, Mg, Sc, Pb c. Pb, Sc, Mg, P,
C d. -4, 2, 2 and 3, -3, 2 and 4
85
Practice Problems Answers
2. e. C 2, 1, 0, 1/2 Mg 3, 0, 0, -1/2 Sc
3, 2, 2, 1/2 P 3,1,-1,1/2 Pb 6, 1, 0,
1/2 3. N, Pd, Cr
86
Electrons and Quantum Numbers
  • n l ml ms
  • 1 0 0 1/2
  • 1 0 0 -1/2

87
Electrons and Quantum Numbers
n l ml ms 2 0
0 1/2 2 0 0 -1/2
  • 2 1 1 1/2
  • 2 1 0 1/2
  • 2 1 -1 1/2
  • 2 1 1 -1/2
  • 2 1 0 -1/2
  • 2 1 -1 -1/2

88
Hydrogen
  • Group 1A
  • Atomic number 1
  • 1 electron

1s1
89
Helium
  • Group 8A
  • Atomic number 2
  • 2 electrons

1s
1s2
90
Lithium
  • Group 1A
  • Atomic number 3
  • 3 electrons

1s
2s
1s22s1
91
Beryllium
  • Group 2A
  • Atomic number 4
  • 4 electrons

1s
2s
1s22s2
92
Boron
  • Group 3A
  • Atomic number 5
  • 5 electrons

1s
2s
2p

1s22s22p1
93
Carbon
  • Group 4A
  • Atomic number 6
  • 6 electrons

1s
2s
2p

1s22s22p2
94
Nitrogen
  • Group 5A
  • Atomic number 7
  • 7 electrons

1s
2s
2p

1s22s22p3
95
Oxygen
  • Group 6A
  • Atomic number 8
  • 8 electrons

1s
2s
2p

1s22s22p4
96
Fluorine
  • Group 7A
  • Atomic number 9
  • 9 electrons

1s
2s
2p

1s22s22p5
97
Neon
  • Group 8A
  • Atomic number 10
  • 10 electrons

1s
2s
2p

1s22s22p6
98
Sample Question
Give the spectroscopic notation for
P. 1s22s22p63s23p3 Determine the set of
4 quantum numbers for the last electron in.
(3, 1, -1, 1/2) Give a box diagram for the
valence electrons.
99
Sodium
  • Group 1A
  • Atomic number 11
  • 11 electrons

2s
2p
3s

1s22s22p63s1
100
Magnesium
  • Group 2A
  • Atomic number 12
  • 12 electrons

2s
2p
3s

1s22s22p63s2
101
Aluminum
  • Group 3A
  • Atomic number 13
  • 13 electrons

2s
3s
3p

1s22s22p63s23p1
102
Silicon
  • Group 4A
  • Atomic number 14
  • 14 electrons

2s
3s
3p

1s22s22p63s23p2
103
Phosphorus
  • Group 5A
  • Atomic number 15
  • 15 electrons

2s
1s22s22p63s23p3
104
Sulfur
  • Group 4A
  • Atomic number 16
  • 16 electrons

2s
3s
3p

1s22s22p63s23p4
105
Chlorine
  • Group 7A
  • Atomic number 17
  • 17 electrons

2s
3s
3p

1s22s22p63s23p5
106
Argon
  • Group 8A
  • Atomic number 18
  • 18 electrons

2s
3s
3p

1s22s22p63s23p6
107
Potassium
  • Group 1A
  • Atomic number 19
  • 19 electrons

2s
3s
3p
4s

1s22s22p63s23p64s1
108
Calcium
  • Group 2A
  • Atomic number 20
  • 20 electrons

2s
3s
4s
1s22s22p63s23p64s2
109
Scandium
  • B Group
  • Atomic number 21
  • 21 electrons

3d

1s22s22p63s23p64s23d1
110
Sample Questions
1. Write the noble gas configuration for a) Ge
Ar 4s23d104p2
b) Cs
Xe 6s1
c) Pb
Xe 6s24f145d106p2
111
Sample Questions
2. What element has (4, 2, 1, -1/2) as its last
electron in? Use a valence box diagram and give
the valence configuration for this element.
4d7 Rh
2 1 0 -1 -2
5s2 or 5s24d7
112
Sample Questions
3. Give the symbol of the element with the lowest
atomic number that a. is a noble gas with no p
electrons. b. is in period 3 and has 2 unpaired
electrons. c. is in period 2 and has 2 filled p
orbitals. d. has 4 electrons in level 2.
He Si F C
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