Orbital Energies
Orbital Energies
4f
Depend on:
Chapter 5c:
4d
• Magnitude of positive
charge on nucleus
(# of protons =
atomic number)
4p
3d
Energy
Electron Energies
and Configurations
4s
• Distance of electrons
from nucleus
(quantum number n)
3p
3s
• Screening of nucleus by
other electrons
(quantum number l:
s, p, d, f shapes)
2p
2s
1s
Chem 111
Dr. Gentry
Nuclear Screening
Electrical-Attraction Energy
Z e2
*C
r
E= −
e–
(electron-electron repulsion)
electron in
2s orbital
r
Z
r
= nuclear charge
= distance of electron
from nucleus
e = electrical charge on
electron
C = a constant
Z+
Zeff = “Effective” nuclear
charge seen by
different electrons
e–
electrons in
1s orbital
e–
e–
Z+
For Li atom (Z = 3)
Zeff (1s) = 2.69
Zeff (2s) = 1.28
• More negative energies have stronger attraction
• Compare Fluorine (Z = 9) to Carbon (Z = 6)
As go across row,
Electrons close to nucleus tend to hide ( or “shield”) some
of the nuclear charge from electrons further away since
electrons repel each other
Z gets bigger = more attraction
• Compare size of Fluorine 2p orbital to Iodine 5p orbital
As go down column, r gets bigger = less attraction
Nuclear Screening
(causes s,p,d,f orbitals to be different energies)
H
Z
1s
Effective Nuclear Charges, Zeff
1
E= −
1.000
Z eff e 2
Zeff = “Effective” nuclear charge seen by different electrons
Z
s
p
d
f
Less
Effective Nuclear Charge
Less shielded
More shielded
Larger Zeff
Less Zeff
1.688
Li
Be
B
C
N
O
F
3
4
5
6
7
8
9
1s
2.691
3.685
2s
1.279
1.912
2p
Probability of Finding Electron Close to Nucleus
More
r
4.680
5.673
He
2
*C
6.665
7.658
8.650
Ne
10
9.642
2.576
3.217
3.847
4.492
5.128
5.758
2.421
3.136
3.834
4.453
5.100
5.758
Na
Mg
Al
Si
P
S
Cl
Z
11
12
13
14
15
16
17
Ar
18
1s
10.626
11.609
12.591
13.575
14.558
15.541
16.524
17.508
2s
6.571
7.392
8.214
9.020
9.825
10.629
11.430
12.230
2p
6.802
7.826
8.963
9.945
10.961
11.977
12.993
14.008
3s
2.507
3.308
3p
4.117
4.903
5.642
6.367
7.068
7.757
4.066
4.285
4.886
5.482
6.116
6.764
Wikopedia
1
E= −
Z eff e
Depend on:
*C
4d
• Magnitude of positive
charge on nucleus
(# of protons =
Z atomic number)
4p
3d
Energy
(start placing electrons in lowest energy orbitals)
4f
2
r
Electron Configuration of Atoms
Orbital Energies
Orbital Energies
4s
2p6
?
3p6
4p6
5p6
• Distance of electrons
from nucleus
(quantum number n)
3p
3s
6p6
• Screening of nucleus by
other electrons
(quantum number l:
s, p, d, f shapes)
2p
2s
1s
Electron Configuration of Atoms
Orbital Energies
And Filling Pattern
4f
4d
• Aufbau Principle: Fill lowest energy first:
Electrons start filling orbitals with lowest energies.
4p
3d
Energy
• Pauli Exclusion Principle:
Two electrons cannot have same 4 quantum numbers (n, l, ml, ms)
⇒ Can only be two electrons in an orbital (e.g. 2px),
one with spin up and one with spin down
4s
3p
3s
• Hund’s Rule:
2p
When filling orbitals in same subshell, maximize number of
parallel spins.
2s
1s
First fill 1 electron each in 2px, 2py, 2pz, each with spin up
Aufbau Principle
(1*s)
(3*p)
(5*d)
(7*f)
# of suborbitals
(2)
(6)
(10)
(14)
# of electrons
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
5f
5d
Orbital Energies
And Filling Pattern
4f
6d
5d
4d
3d
5f
4f
8
(2)
(6)
5s
7s
7p
6s
6p
6d
5s
5p
5d
5f
3p
4s
4p
4d
4f
3s
3s
3p
3d
2s
2p
4p
3d
4s
7
6
2p
5
4
2s
3
2
5p
4d
Energy
Order of filling based
on increasing energies
of each orbital
1s
(10)
(14)
8
7
6
5
4
2
3
1
1s
1
2
↑↓
1s
↑
2s
2px 2py 2pz
↑↓
1s
↑↓
2s
2px 2py 2pz
B
↑↓
1s
↑↓
2s
C
↑↓
1s
O
Ne
Li
Be
Shorthand “Condensed” Notation
Electron Configuration
1s2 2s1
1s2 2s2
Ne ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s 2px 2py 2pz
↑
2px 2py 2pz
1s2 2s2 2p1
S
↑↓
2s
↑
↑
2px 2py 2pz
1s2 2s2 2p2
↑↓
1s
↑↓
2s
↑↓
2px
↑
↑
2py 2pz
1s2 2s2 2p4
↑↓
1s
↑↓
2s
↑↓ ↑↓ ↑↓
2px 2py 2pz
1s2 2s2 2p6
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s 2px 2py 2pz
↑↓
3s
↑↓
3px
↑
↑
3py 3pz
Ne: 1s2 2s2 2p6
S:
1s2 2s2 2p6 3s2 3p4
… OR …
Electron Configuration of Atoms
[Ne] 3s2 3p4
Refer to previous Nobel Gas
Electron Configuration of Atoms
(start placing electrons in lowest energy orbitals)
• Give the ground-state electron configurations for:
Ne (Z = 10)
Mn (Z = 25)
Rb (Z = 37)
Eu (Z = 63)
Hg (Z = 80)
Am (Z = 95)
2p6
3p6
4p6
5p6
• Identify elements with ground-state configurations:
6p6
1s2 2s2 2p4
1s2 2s2 2p6
[Ar] 4s2 3d1
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d6
[Xe] 6s2 4f14 5d10 6p5
Electron Configurations
(Figure 5.16)
The “Rule Breakers”
(anomalous configurations)
Anomalous Electron Configurations
• Result from unusual stability of half-filled & fully-filled
subshells.
Chromium should be [Ar] 4s 2 3d 4, but is [Ar] 4s 1 3d 5
Copper
should be [Ar] 4s 2 3d 9, but is [Ar] 4s 1 3d 10
• In the second transition series this is even more
pronounced
Nb, Mo, Ru, Rh, Pd, and Ag have anomalous
configurations
See Figure 5.16 in text
3
Ionization Energy
Across horizontal row (e.g., Na Cl ),
stay in same shell (n=3), but nucleus
is more positive so pulls electrons closer
Atomic Number
•
•
First Ionization Energy, kJ/mol
Atomic Radius, 10-12m
Atomic Radius
Energy needed to remove an electron
Depends on nuclear charge (Z) and distance (n & l)
Efree elec. = 0
Ion.
Energy
Eorbital = −
Z e2
*C
r
Atomic Number
• High electron affinity
(want to add electrons)
• Nonmetallic
• High ionization potential
• Small radius (size)
(electrons pulled in tight)
• Low electron affinity
(willing to give up electrons)
• Metallic
• Low ionization potential
• Large radius (size)
(floppy, loose e- cloud)
4