EMR and the atom: Part Deux
Electron Configurations
http://imagers.gsfc.nasa.gov/ems/waves3.html
First, a review…
1. What do scientists mean when they say light
is quantized?
2. Compare a ground state and an excited state
electron.
3. Explain the origin of the atomic emission
spectrum of an element.
4. What do we know so far about the
arrangement of electrons in atoms?
What you’ve seen so far….
Model of a
Nitrogen (z=7)
atom
Which is really not true- why?
• Because orbitals- the “electron
cloud” are
–3-D, not flat
–are not round in most cases
–e-s are spread out as much as
possible
–(e-s are moving very rapidly)
Orbitals
• The electrons are spread out in orbitals that have
varying
– Shapes
– Energy (distance from nucleus)
• The orbitals are described in regards to their
quantum numbers
– Descriptions that are descriptive and hierarchical
– There are 4 numbers that describe an orbital
• Written as follows: (#, #, #, ±#)
Principal quantum number (n)
The first number (1, #, #,±#)
• Describe the
– distance from the nucleus of the orbital
– The energy of the orbital
• Values for n are integers
– The smallest possible value is 1
• As the distance from the nucleus (and therefore
energy) increases, the number increases
Quantum numbers
There periodic table and n
• The 7 periods on the periodic table
correspond to n values
• Each period has a unique n value
– For the 1st period, n=1
– For the 2nd period, n=2
– And so on….
The shape of things
• Is the shape of the orbital
• There are 4 shapes (although we only deal with the
first three)
–
–
–
–
s = sphere
p = peanut
d = double peanut
f = flower
The s orbital
http://www.sfu.ca/~nbranda/28xweb/images/s_orbital.gif
p orbitals
d orbitals
d orbitals
f orbitals
General tutorials for electron configuration stuff
• some slides in this PowerPoint are from this site
already
• http://www.wwnorton.com/college/chemistry
/gilbert/tutorials/ch3.htm
• See key equations and concepts (select from menu
on the left), as well as the looking through the
overview where to the tutorials are listed (links for
just those are on the left, too)
Orbitals
• Denote the orbital sublevel that is filled
• It is the third number in the description (#,#,1, ±#)
– s orbitals have one sublevel; a sphere has one orientation
in space
– p orbitals have three sublevels; 3 orientations in space
– d orbitals have five sublevels; 5 orientations in space
– f orbitals have seven sublevels; 7 orientations in space
“Flavors” of ml
• s orbitals have
one sublevel; a
sphere has one
orientation in
space
“Flavors” of ml
• p orbitals have three sublevels; 3 orientations in
space
“Flavors” of ml
• d orbitals have
five sublevels; 5
orientations in
space
“Flavors” of ml
• f orbitals have
seven
sublevels; 7
orientations in
space
Spin
• It is the last number in the description
(#,#,#,±½)
• Spin is +½ or -½
– Up or down
How we use this….
• There is a specific order to how the e- fill the
orbitals; it is not random
– Although there are exceptions to the rules (last
thing we do)
The principles of e- configuration
• The Aufbau (next) Principle:
– That e- fill the lowest energy sublevel before going to the
next sublevel
• The Pauli Exclusion Principle:
– That e-s are paired according to opposite spins
• Hund’s Rule:
– e-s spread out in equal energy sublevels before placing
electrons
• The first level to fill is the 1s level
– It is the lowest energy sublevel
– It holds two electrons
• They are oppositely paired (up and down- ↑↓)
• Each sublevel (each __) holds 2 electrons
Next…
• The second sublevel is the 2s sublevel
• It also holds 2 electrons (because s holds 2,
not because of the number),
• also oppositely paired ↑↓
1s2, 2s2,then comes 2p6
• So, as it states above
– 1s fills, 2s fills ,then
comes 2p
• It holds up to six
electrons
• Because p orbitals hold
6 electrons
Next…
• From 2p,
–
–
–
–
3s fills with 2e-, then onto
3p, with 6e- then
4s with 2e- followed by
3d with 10e- (because d holds
10e-)
– Then 4p with 6e-
• Notice, you follow the
arrows
• Remember, the number of
electrons comes from the
letter (the orbital’s
momentum,ml)
• The sublevels of the
orbitals are first filled,
then you continue onto
the next level (Aufbau)
• Also be sure to place one
electron in each sublevel
prior to filling the level
(↑ ↑ ↑ and not ↑↓
↑ _) (Hund)
• e-s must be paired with es of opposite spin (↑↓,
not ↑↑ or ↓↓) (Pauli)
Putting it all together…
• Carbon (neutral, so 6 electrons)
• What this would look like:
↑↓ ↑↓ ↑ ↑ _
1s 2s 2p
(notice there are 6 arrows for 6 electrons)
• This can also be written as 1s2 2s2 2p2
• Notice the superscripts add up to 6
There are some exceptions…
• This is because some energy levels are very close together
– electrons are able to move between close orbitals in order to minimize
repulsion
• Example: the 4s and 3d orbitals are very close in energy
• So exceptions for some period 4 d block elements occur
– Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4
– Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5
– Because it takes less energy to split the electrons between the 5
sublevels than it does to put them together in the 4s and 3d