Download Chapter 8: Electron Configurations and Periodicity

Electron Spin & the Pauli Exclusion Principle
Chapter 8:
Electron Configurations
and Periodicity
4. spin quantum number, ms
◆
◆
electrons behave like they are
spinning around an axis
spinning of a charged particle
creates a magnetic field
magnetic field created has a direction
if e– spins clockwise, magnetic field is one
direction (ms = +!)
if e– spins counter-clockwise, magnetic field is
opposite direction (ms = –!)
3 quantum numbers (n, l, ml) define the energy,
size, shape, and spatial orientation of each atomic
orbital.
To explain how electrons populate atomic orbitals
in an atom, we need a 4th quantum number.
Experimental Evidence of Electron Spin:
Stern & Gerlach (1921)
How Do Electrons Populate Atomic Orbitals?
Pauli Exclusion Principle (1925)
No 2 electrons in an atom can have the same set
of 4 quantum numbers.
◆
n, l, ml define the atomic orbital
◆
ms will define the electrons in the orbital
If there are only 2 possible ms values, then each
atomic orbital can hold no more than 2 electrons;
◆
Effective Nuclear Charge (Zeff)
What is the charge felt by an electron in an atom?
◆
depend on n and l of the orbital where the e– lives
◆
the farther the electron is from the nucleus:
the less the force of attraction to the nucleus
the greater the e– - e– repulsion
◆
an outer shell electron is “shielded” by inner shell e–’s
specifically, one e– must have ms = +!
and the other e– must have ms = –!
These 2 e–’s in an atomic orbital are said to be
spin-paired.
Effective Nuclear Charge (Zeff)
Zeff = Zactual – shielding factor
Zeff and Atomic Orbitals
◆
relationship between Zeff and n
probability of an e– being close to the nucleus:
n=1>n=2>n=3
Zeff felt by an e– in an orbital:
n=1>n=2>n=3
energy of orbital:
n=1<n=2<n=3
Electron Configurations
Zeff and Atomic Orbitals
◆
relationship between Zeff and l (within a same shell)
◆
probability of an e– being close to the nucleus:
l=0>l=1>l=2
or:
s>p>d
Zeff felt by an e– in an orbital:
l=0>l=1>l=2
or:
s>p>d
energy of orbital:
s<p<d
The Aufbau Principle
1. Lower energy orbitals fill before higher energy orbitals.
◆ use the Relative Energies of Atomic Orbitals diagram
2. An orbital can accommodate a maximum of 2 e–’s
which must be spin-paired.
◆ Pauli Exclusion Principle
3. Hund’s Rule:
If 2 or more degenerate orbitals are available, one e–
will go into each orbital until all are half-full.
◆
the e–’s in the singularly populated orbitals must
have the same ms
After all degenerate orbitals are half-full, then a 2nd
e– may be added to fill the orbitals.
◆
give complete electronic description for every element
◊
predict orbitals occupied by electrons
◊
write electron configurations
◊
draw orbital diagrams
follow set of 3 rules: Aufbau (“building up”) Principle
◊
each successive electron will go into the lowest
energy orbital available
◊
this results in the lowest energy, ground state
configuration
In What Order Do Atomic Orbitals Populate?
Lower Energy to Higher Energy
examples:
Write the ground state electron configurations, and
complete an orbital diagram for neutral atoms of the
following elements.
N (7 e–) :
1s2
Al (13 e–) :
↑↓
1s
2s2 2p3
↑
↑↓
↑↓
1s
2s
2p
1s2
2s2 2p6
3s2 3p1
↑↓
2s
↑↓
↑↓
↑
↑
↑↓
↑↓
2p
3s
Sc (21 electrons):
↑↓
↑↓
1s
2s
↑↓
↑↓
↑↓
2p
↑
↑↓
3p
outer shell electrons
core electrons
valence electrons
paramagnetic
diamagnetic
4s
1s2 2s2 2p6 3s2 3p6 4s2 3d1
↑↓
↑↓
3s
Ga (31 electrons):
Some Terminology:
inner shell electrons
Closed Shells and Subshells & Using Noble Gas Core
Symbolism in Electron Configurations
↑↓
↑↓
↑↓
3p
↑↓
↑
4s
3d
[Ar] 4s2 3d10 4p1
↑↓
↑↓
3d
↑↓
↑↓
↑
4p
Some Anomalies in Electron Configurations:
◆
result from unusual stability of half-filled or
completely filled shells or subshells
ex: Cr predict:
↑↓
↑
[Ar] 4s2 3d4
↑
↑
4s
3d
ex: Cu predict:
↑↓
4s
◆
!
↑
↑↓
↑↓
↑↓
3d
actual:
↑
↑
[Ar] 4s1 3d5
↑
4s
[Ar] 4s2 3d9
↑↓
↑
!
4s
↑
↑
3d
actual:
↑
↑
↑↓
[Ar] 4s1 3d10
↑↓
↑↓
↑↓
↑↓
3d
heavy elements (above atomic # 40) ∆E between
orbitals is smaller, so anomalies are common
Electron Configurations and the Periodic Table
Periodic Trends
Periodic Trend in Zeff
increase
The goal is to use our understanding of electron
configurations and Zeff to understand trends in:
◆
atomic radius
◆
ionization energy
◆
electron affinity
How does a given property change from
left to right across periodic table?
How does a given property change from
top to bottom of periodic table?
Atomic Radius
radius (typically in pm or Å) of neutral atoms of elements
trend:
atomic radius decreases left to right across the
periodic table
atomic radius increases top to bottom of the
periodic table
decrease
i
n
c
r
e
a
s
e
Atomic Radius
d
e
c
r
e
a
s
e
Effective Nuclear Charge,
Zeff
Na
Mg
e–
configuration
[Ne]3s1
[Ne]3s2
Al
actual nuclear
charge
11
12
13
14
15
16
17
18
# core e–’s
10
10
10
10
10
10
10
10
# valence e–’s
1
2
3
4
5
6
7
8
Zeff
+1
+2
+3
+4
+5
+6
+7
+8
[Ne]3s2
Si
3p1
[Ne]3s2
P
3p2
[Ne]3s2
S
3p3
[Ne]3s2
Cl
3p4
Periodic Trend in Atomic Radius
[Ne]3s2
Ar
3p5
[Ne]3s2 3p6
Ionization Energy
Periodic Trend in Atomic Radius
ionization energy – the energy required to remove
an electron from a gas phase
atom or ion in its ground state
X (g) " X+ (g) + e– ;
d
e
c
r
e
a
s
e
Periodic Trend in Ionization Energy
endothermic
increase
Ionization Energy
Periodic Trend in Ionization Energy
consider successive ionization energies:
M (g) ! M+ (g) + e–
M+ (g) ! M2+ (g) + e–
M2+ (g) ! M3+ (g) + e–
◆
Note the unexpected changes between groups IIA & IIIA,
and groups VA & VIA. Why? Think about e– configurations.
◆
1st ionization energy
2nd ionization energy
3rd ionization energy
It becomes successively harder to remove an e– from a positively
charged species because of forces of electrostatic attraction.
Periodic Trend in Ionization Energy
◆
removing a core e– costs MUCH more energy than
removing a valence e–
Electron Affinity
electron affinity – change in energy that occurs when an
electron is added to an isolated gas phase atom.
X (g) + e– ! X– (g)
*
d
e
c
r
e
a
s
e
◆
Valence electrons are most easily lost during ionization,
and are gained, lost, or shared during chemical reactions.
* increase
Electron Affinity
* increase means becomes larger, negative value ∴ more
favorable for anion formation;
decrease means becomes smaller, negative value ∴ less
energy released and less favorable for anion formation
Periodic Trend in Electron Affinity
note: where electron affinity values are > 0, anion formation
is very unfavorable; alkaline earth metals & the noble gases
Periodic Trend in Electron Affinity