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Electron Configurations: In what order do electrons occupy available orbitals?
Chapter Outline
Orbital Energy Levels for Hydrogen Atoms
3.1 Waves of Light
3.2 Atomic Spectra
3.3 Particles of Light: Quantum Theory
3.4 The Hydrogen Spectrum and the Bohr Model
3.5 Electrons as Waves
3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals
3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions
3.10 The Sizes of Atoms and Ions
3.11 Ionization Energies
3.12 Electron Affinities
E
3s
3p
2s
2p
3d
1s
1
Energy of orbitals in multi-electron atoms
1s
"Penetration"
3+2=5
4+0=4
3+1=4
3+0=3
2+1=3
2+0=2
1+0=1
Probability of finding the electron →
Energy depends on n + l
s>p>d>f
2s
3s
4s
2p
3p
4p
3d
4f
for the same shell (e.g.
n=4) the s-electron
penetrates closer to
the nucleus and feels a
stronger nuclear pull or
charge.
4d
distance from nucleus →
http://www.pha.jhu.edu/~rt19/hydro/img73.gif
Aufbau Principle - the lowest energy orbitals fill up first
Filling order of orbitals in multi-electron atoms
Shorthand description of orbital occupancy
1. No more than 2 electrons maximum per orbital
2. Electrons occupy orbitals in such a way to minimize the
total energy of the atom = “Aufbau Principle”
(use filling order diagram)
3. No 2 electrons can have the same 4 quantum numbers
= “Pauli Exclusion Principle” (pair electron spins)
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
ms = +1/2
spin “up”
ms = -1/2
spin
“down”
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•No two electrons in an atom can have the same set of four
quantum numbers (n, l, ml, ms)
4. When filling a subshell, electrons occupy empty
orbitals first before pairing up = “Hund’s Rule”
•electrons must "pair up" before entering the same orbital







Px
Py
Pz
NOT


Px
Py
Pz
“orbital box diagram”
Electron Shells and Orbitals
• Orbitals that have the exact same energy level are called degenerate.
• Core electrons are those in the filled, inner shells in an atom and are not involved in chemical reactions.



Px
Py
Pz
+
• Valence electrons are those in the outermost shell of an atom and have the most influence on the atom’s chemical behavior.
H:
He:
Li:
Be:
B:
C:
N:
O:
F:
Ne:
Electron Configurations from the Periodic Table
Transition metals are characterized by having incompletely
filled d-subshells (or form cations as such).
n‐1
n‐2
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Electron Configurations: Ions
Chapter Outline
• Formation of Ions:
3.1 Waves of Light
3.2 Atomic Spectra
3.3 Particles of Light: Quantum Theory
3.4 The Hydrogen Spectrum and the Bohr Model
3.5 Electrons as Waves
3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals
3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions
3.10 The Sizes of Atoms and Ions
3.11 Ionization Energies
3.12 Electron Affinities
– Gain/loss of valence electrons to achieve stable electron configuration (filled shell = “octet rule”).
– Cations:
– Anions:
– Isoelectronic:
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Cations of Transition Metals
Sample Exercise 3.11:
Determining Isoelectronic Species in Main Group Ions
Fe
a) Determine the electron configuration of each of the following ions: Mg2+, Cl‐, Ca2+, and O2‐
Cu
b) Which ions are isoelectronic with Ne?
Sn
Pb
Periodic Trends – trends in atomic and ionic radii, ionization energies, and electron affinities
Chapter Outline
3.1 Waves of Light
3.2 Atomic Spectra
3.3 Particles of Light: Quantum Theory
3.4 The Hydrogen Spectrum and the Bohr Model
3.5 Electrons as Waves
3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals
3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions
3.10 The Sizes of Atoms and Ions
3.11 Ionization Energies
3.12 Electron Affinities
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Effective nuclear charge (Zeff) – Inner shell electrons
“SHIELD” the outer shell electrons from the nucleus
Zeff = Z ‐ 
 = shielding constant) Effective nuclear charge (Zeff) – Inner shell electrons
“SHIELD” the outer shell electrons from the nucleus
Down a family -
Zeff  Z – number of inner or core electrons
Across a period -
Z
Core Zeff
Radius (nm)
Na
11
10
1
186
Mg
12
10
2
160
Al
13
10
3
143
Si
14
10
4
132
Trends in Effective Nuclear Charge (Zeff)
and the Shielding Effect
Core Zeff
Z
Radius (nm)
Na
11
10
1
186
K
19
18
1
227
Rb
37
36
1
247
Cs
55
54
1
265
Atomic, Metallic, Ionic Radii
increasing Shielding
increasing Zeff
For diatomic molecules, equal to covalent radius (one‐half the distance between nuclei).
Trends in Atomic Size for the “Representative (Main Group) Elements”
For metals, equal to metallic radius (one‐
half the distance between nuclei in metal lattice).
For ions, ionic radius equals one‐half the distance between ions in ionic crystal lattice.
Radius of Ions
Increasing Atomic Size
Decreasing Atomic Size
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
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Sample Exercise 3.13:
Ordering Atoms and Ions by Size
Radii of Atoms and Ions
must compare cations to cations and anions to anions
Decreasing Ionic Radius
Arrange each by size from largest to smallest:
Increasing Ionic Radius
(a) O, P, S
(b) Na+, Na, K
Ionization energy is the minimum energy (kJ/mol)
required to remove an electron from a gaseous
atom in its ground state.
Chapter Outline
3.1 Waves of Light
3.2 Atomic Spectra
3.3 Particles of Light: Quantum Theory
3.4 The Hydrogen Spectrum and the Bohr Model
3.5 Electrons as Waves
3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals
3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions
3.10 The Sizes of Atoms and Ions
3.11 Ionization Energies
3.12 Electron Affinities
I1 + X (g)
X+(g) + e-
I1 first ionization energy
I2 + X (g)
X2+(g) + e-
I2 second ionization energy
I3 + X (g)
X3+g) + e-
I3 third ionization energy
I1 < I2 < I3 < …..
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General Trend in First Ionization Energies
Ionization Energies
Decreasing First Ionization Energy
Increasing First Ionization Energy
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Successive Ionization Energies (kJ/mol)
Sample Exercise 3.14:
Recognizing Trends in Ionization Energies
Arrange Ar, Mg, and P in order of increasing IE
Trends in the 1st Ionization Energy for the 2nd Row
Li
520 kJ/mol 1s22s1  1s2
Be
899
1s22s2  1s22s1
B
801
1s22s22p1  1s22s2
C
1086
1s22s22p2  1s22s22p1
N
1402
1s22s22p3  1s22s22p2
O
1314
1s22s22p4  1s22s22p3
F
1681
1s22s22p5  1s22s22p4
Ne
2081
1s22s22p6  1s22s22p5
Chapter Outline
3.1 Waves of Light
3.2 Atomic Spectra
3.3 Particles of Light: Quantum Theory
3.4 The Hydrogen Spectrum and the Bohr Model
3.5 Electrons as Waves
3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals
3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions
3.10 The Sizes of Atoms and Ions
3.11 Ionization Energies
3.12 Electron Affinities
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Electron affinity is the energy release that occurs
when an electron is accepted by an atom in the gaseous
state to form an anion.
X (g) + e- → X-(g)
Energy released = E.A. (kJ/mol)
F (g) + e- → F-(g)
EA = -328 kJ/mol
O (g)
+ e- → O-
(g)
Periodic Trends in Electron Affinity
EA = -141 kJ/mol
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