Transcript Lecture 21 revised (Slides) October 12

Electron Configurations
• Electrons can be distributed amongst the
subshells/orbitals of an atom in different ways
–producing different electron configurations.
The most stable configuration has the lowest
energy – corresponding to the situation where
electrons get as close to the nucleus as possible
while staying as far away from each other as
possible.
Electron Configurations
• Aufbau process
– Electrons occupy orbitals in a way that
minimizes the energy of the atom.
• Pauli exclusion principle
– No two electrons can have all four quantum
numbers alike.
• Hund’s rule
• When orbitals of identical energy (degenerate
orbitals) are available, electrons initially occupy
these orbitals singly.
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General Chemistry: Chapter 8
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Populating Orbitals
• If we continued with the previous exercise
we’d see that for each value of n there is one s
orbital (unique ml value), for n ≥ 2 three p
orbitals (three ml values) and, for n ≥ 3, five d
orbitals (five ml values: -2, -1, 0, +1, +2). This
means that s, p, d subshells can contain (at
most) 2, 6 and 10 electrons respectively.
Relative subshell energies comes from
experiment – next slide.
FIGURE -37
The order of filling of electronic subshells
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Electron Configurations - Monatomic
Species
• We will usually write condensed electron
configurations which show the number of
electrons per subshell (usually for the ground
state or lowest energy configuration possible).
Excited states or higher energy configurations
will also be considered. We will also write
expanded electron configurations where –
particularly for valence electrons – the number
of electrons per orbital is shown.
“Condensed” Electron Configurations
Example of C Atom – 6 Electrons
Subshell
1s
2s
2p
3s
3p
Maximum e-s
Per Subshell
2
2
6
2
6
Actual # e-s
Per Subshell
2
2
2
0
0
Number of e-s
Accommodated
2
4
6
0
0
Number of e-s
“Left Over”
4
2
0
0
0
Electron Configurations –
“Exponential” Notation
• The information contained in the previous
slide can be represented more compactly using
a notation where “exponents” are used to
indicate the number of electrons per subshell.
For the neutral C atom in its ground electronic
state the electron configuration written in
exponential notation is 1s22s22p2.
“Expanded” Electron Configurations
• A more detailed picture of electron the
electronic “structure” of an atom or monatomic
ion is obtained using so-called “expanded”
configurations. The expanded configurations
reflect Hund’s Rule (based on experiment) –
electrons occupy equivalent orbitals to the
maximum extent possible and with their spins
parallel.
“Expanded” Electron Configurations
Example of C Atom – 6 Electrons
Subshell
1s
2s
2px
2py
2pz
Maximum e-s Actual # e-s
Per Subshell Per Subshell
2
2
2
2
2
2
2
1
1
0
Number of e-s
Accommodated
2
4
5
6
0
Number of e-s
“Left Over”
4
2
1
0
0
“Expanded” Electron Configurations
• For a C atom in its ground state the expanded
electron configuration could be written as
1s22s22px12py1 OR 1s22s22px12pz1 OR
1s22s22py12pz1. All configurations have the
same energy. Expanded configurations help us
identify any unpaired electrons that might be
present. Unpaired electrons are important since
they can give rise to important magnetic
properties (paramagnetism).
Expanded Electron Configurations &
Orbital Diagrams
• The detailed electron configurations of atoms
and monatomic ions can also be represented
using orbital diagrams. Here a box represents
each orbital and arrows indicate both the
“spins” of electrons (ms value +ve or –ve) and
whether the orbital contains 0, 1 or 2 electrons.
The next slide represents the orbital diagram
for the C atom. Note the “parallel spins” for
the two 2p electrons.
Representing Electron Configurations
spdf notation (condensed)
1s22s22p2
spdf notation (expanded)
1s22s22px12py1
spdf notation
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General Chemistry: Chapter 8
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The Aufbau process
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Transition Metals – Electron
Configurations
• Transition metals have surprisingly rich
chemistry. The electronic configurations of the
neutral atoms are relatively complex since d
subshells (with 5 orbitals) are being filled. For
the first series of transition metals the 4s and
3d subshells have similar energies and
surprises are seen for their electronic
configurations. Cr and Cu do not have the
“expected” electron configurations.
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The Aufbau Process – Sc through Zn
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Transition Metal Configurations
• The unexpected electron configurations found
experimentally for Cu and Cr are often
rationalized in terms of a special stability (low
energy) associated with a half full (3d5) and
full (3d10) d subshell. Similar issues arise with
transition metal ions. The large numbers of
unpaired electrons seen for transition metals
gives them interesting magnetic properties.
Transition metal compounds are often
colourful – discussed in higher level courses.
Electron Configurations and the
Periodic Table
FIGURE 8-38
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Valence Shell Configurations
• The occupied shell with the highest value of n
is called the valence shell. When atoms
undergo chemical change electrons in the
valence shell can be lost or shared with other
atoms. The valence shell can also pick up
electrons. Atoms with similar chemical
properties often have the “same” valence shell
electron configuration. For example, Li, Na, K,
Rb, Cs and Fr have an ns1 valence shell
configuration.
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Electron Configurations - Examples
• 1. Write condensed electron configurations
representing the ground electronic states of the
P atom and the P3- ion.
• 2. Write condensed electron configurations
representing the ground electronic states of the
Sr atom and the Sr2+ ion.
• 3. Write condensed electron configurations for
the ground and two excited electronic states of
the Na atom.
Electron Configurations - Examples
• 4. Construct orbital diagrams for (a) the Al
atom, (b) the Si atom (for Bill Gates), (c) the S
atom and (d) the S2- ion.
• 5. How many unpaired electrons are there in a
neutral arsenic (As) atom?
• 6. Write a set of four possible quantum number
values (n, l, ml and ms) for an electron in the
(a) valence shell of Mg and (b) the highest
occupied orbital (energy) of Pb.
Electron Configurations - Examples
• 7. Write chemical symbols for five monatomic
species having the ground state electron
configuration 1s22s22p63s23p6. How many
electrons are protons are contained in each
species?
• 8. How many unpaired electrons do neutral Al,
Si, P, S, Cl and Ar atoms possess?
Final Note on Nodes
• The H atom wave functions tell us that there
are only radial nodes for the 1s, 2s, 3s..
orbitals. Angular or planar nodes become
important as we move to p orbitals (one
planar node) and d orbitals (2 planar nodes).
This is seen in the slide on the next table and
graphical representations of d orbitals.
(Orbital shapes impt in chemical bonding.)
Hydrogen Atom Wavefunctions –
Number of Nodes
Orbital
Designation
Total # Nodes
(n-1)
Planar Nodes
(l)
Radial Nodes
1s
0
0 (s orbital)
0
2s
1
0 (s orbital)
1
2p
1
1 (p orbital)
0
3s
2
0 (s orbital)
2
3p
2
1 (p orbital)
1
3d
2
2 (d orbital)
0
FIGURE 8-30
Representations of the five d orbitals
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The Periodic Table
• In studying the electronic structure of atoms
we mentioned that chemical families of
elements have similar valence shell electron
configurations. Historically, however, a
detailed knowledge of atomic structure came
subsequent to the observation that groups of
elements had similar chemical properties.
The Periodic Table
• In the modern Periodic Table elements are
arranged in order of increasing atomic number
so that groups of elements with similar
chemical properties appear in columns. The
100+ known elements are commonly divided
into three groups – the metals, non-metals and
the metalloids. These three sets of elements
have rather different physical properties.
Metals and Nonmetals and Their Ions
• Metals
– Good conductors of heat and electricity.
– Malleable and ductile.
– Moderate to high melting points.
• Nonmetals
– Nonconductors of heat and electricity.
– Brittle solids.
– Some are gases at room temperature.
• Metalloids
– Metallic and non-metallic properties
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