Transcript Electron Configurations PPT

Electron Configurations
• Of the three major subatomic particles, the
electron plays the most significant role in
determining the physical and chemical
properties of an element. The arrangement
of elements in the periodic table depends
on these properties.
• Thus, there should be some relationship
between the electron configurations of the
elements and their placement in the table.
Electron Configurations
• The orbital names s, p, d, and f stand for
names given to groups of lines in the
spectra of the alkali metals. These line
groups are called sharp, principal, diffuse,
and fundamental.
Electron Configurations
s
p
1
2
3
4
5
6
7
d (n-1)
f (n-2)
6
7
Electron Configurations
• The electron configuration of an atom
denotes the distribution of electrons
among available shells.
• The standard notation lists the subshell
symbols, one after another.
• The number of electrons contained in each
subshell is stated explicitly.
• For example, the electron configuration of
beryllium, with an atomic (and electron)
number of 4, is 1s22s2 or [He]2s2.
Electron Configurations
• C: 1s2 2s2 2p2
• Ne:1s2 2s2 2p6
• S: 1s2 2s2 2p6 3s2 3p4
or [He] 2s2 2p2
or [He] 2s2 2p6
or [Ne]3s2 3p4
s
p
1
2
3
4
5
6
7
d (n-1)
f (n-2)
6
7
Aufbau Principle
• Electrons fill orbitals starting at the lowest
available energy states before filling higher
states (e.g. 1s before 2s).
• The number of electrons that can occupy
each orbital is limited by the Pauli
Exclusion Principle (each orbital can
hold two electrons with opposite spins).
The rules of the Aufbau Principle
are:
“The Lazy Tenant Rule”
• 1.Electrons are placed in the lowest
energetically available subshell.
• 2. An orbital can hold at most 2 electrons.
• 3.If two or more energetically equivalent
orbitals are available (e.g., p, d etc.) then
electrons should be spread out before they
are paired up (Hund's rule).
Hund's Rule
“The Empty Bus Seat Rule”
• If multiple orbitals of the same energy are
available, Hund's Rule says that
unoccupied orbitals will be filled before
occupied orbitals are reused (by electrons
having different spins).
WRONG
RIGHT
Increasing energy
7p
7s
6p
6s
5p
5s
4d
3d
3p
3s
2p
1s
5d
4p
4s
2s
6d
5f
4f
The Noble Gases
• These are the elements in which the outermost s
and p subshells are filled. The noble gases have
full outer shells; notice that these elements have
filled outermost s and p sublevels
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Helium
Neon
Argon
Krypton
1s2
1s22s22p6
1s22s22p63s23p6
[He]
[Ne]
[Ar]
1s22s22p63s23p63d104s24p6
[Kr]
Representative Elements
• In these elements, the outermost s or p
sublevel is only partially filled. There are
three groups of representative elements:
• Group 1
alkali metals
• Group 2
alkaline earth metals
• Group 7
halogens
Representative Elements
• For any representative element, the group
number equals the number of electrons in the
outermost energy level (valence electrons)
• Potassium
1s22s22p63s23p64s1
• Carbon, silicon, and germanium, in Group 4,
have four electrons in the outermost energy
level
• Carbon
1s22s22p2
• Silicon
1s22s22p63s23p2
• Germanium
1s22s22p63s23p63d104s24p2
Transition Metals
• These are metallic elements in which the
outermost s sublevel and nearby d
sublevel contain electrons. The transition
elements are characterized by addition of
electrons to the d orbitals
Inner Transition Metals
• These are the metallic elements in which
the outermost s sublevel and nearby f
sublevel generally contain electrons
Write the electronic configurations
for the following elements
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O
Na
Ar
Fe
Ca
Ce
1s22s22p4 or [He]2s22p4
1s22s22p63s1 or [Ne]3s1
1s22s22p63s23p6 or [Ne]3s23p6
1s22s22p63s23p63d64s2 or [Ar]3d64s2
1s22s22p63s23p64s2 or [Ar]4s2
1s22s22p63s23p63d104s24p64d105s25p64f15d16s2
or [Xe]4f15d16s2
Half-Full and Full Subshells
• full subshell: fully-filled shells are lower in
energy than partially-filled shells (i.e.
Noble Gases)
• half-filled subshells: lower in energy than
partially-filled subshells
• Cu exception: [Ar] 4s13d10 rather than [Ar]
4s23d9
• Cr exception: [Ar]4s13d5 rather than
[Ar]4s23d4