Electron
Configuration
Writing e- configurations
Drawing orbital notations
Quantum Mechanics Model• describes the electron cloud around the
nucleus as a 3-D wave in which the
electrons move. There are different energy
levels in this cloud with limits the number of
electrons in each level.
Heisenberg Uncertainty Principle• states that it is impossible to know both the
velocity and position of a moving body (such
as the electron or photon) at any one time.
Energy levels• Referred to as principal energy levels
(denoted by the letter “n”).
• Total of 7 energy levels (look at the periodic
table…there are 7 rows called PERIODS or
SERIES)
Energy levels• Lowest energy level, n=1, is closest to the
nucleus and has a low amount of electron
energy.
• The electron energy is called quanta.
• As the value of “n” increases farther from the
nucleus, the energy also increases.
Sublevels• Each energy level is broken into sublevels
(denoted by the letter “l”).
• There are four available sublevels- s, p, d,
and f.
Sublevels•
•
•
•
“s” can hold up to 2 electrons
“p” can hold up to 6 electrons
“d” can hold up to 10 electrons
“f” can hold up to 14 electrons
Electron Configuration• Is the arrangement of electrons in an atom.
[It is like writing an address for the electrons
in their energy levels.]
Note: Remember the number of electrons that
are being counted come from the atomic
number…if the atom is neutral, the
Protons = electrons
Energy Diagram
Rules to using the Energy Diagram :
1. Always start at the lowest energy level (1s).
This is known as the Aufbau Principle.
2. Follow the diagonal arrows after leaving 1s.
3. You must fill up a sublevel before moving
on to the next available sublevel.
• Guided Practice: Write out the electron
configuration for magnesium, phosphorus,
zinc, and bromine.
Drawing orbital notation• Follow the energy diagram to write out the
electron configuration.
• Under the configuration, draw the boxes to
represent the orbitals.
• Use boxes to represent an orbital (where the
electrons are found within the sublevels.)
Drawing orbital notation• Only 2 electrons can fit in an orbital and they
must having opposite spins (called the Pauli
Exclusion Principle.)
• Electrons will be represented with arrows, ↑
and ↓ .
• “s” can hold up to 2 electrons, it will have
1 box
• “p” can hold up to 6 electrons, it will have
3 boxes
• “d” can hold up to 10 electrons, it will
have 5 boxes
• “f” can hold up to 14 electrons, it will
have 7 boxes
• Hund’s Rule states that equal energy
electrons must be spread out evenly in the
orbital (boxes). In other words, place all up
arrows in the boxes before filling in with the
down arrows. See pg.136.
• Guided Practice: Using the electron
configuration for Mg, P, Zn, and Br, draw the
orbital notation for each.
Valence Electrons
• The outermost electrons;
• involved in chemical bonding (we will use
these in later chapters to form ions and form
bonds);
• represented by the electrons found in the “s”
and “p” sublevels.
Valence Electrons
• After writing the electron configuration out,
pick the highest energy level you have
written…
• The electrons found in these levels will be
either “s” or “s and p” electrons…
• These are the valence electrons…never
exceeding a maximum of 8…(octet rule).
Valence Electrons
• A quick way to tell the valence electrons of
the main group elements (the tall columns) is
to look at the roman numeral of the group…
Valence Electrons
Ex.
I A sodium….1 valence electron
II A magnesium…2 valence electrons
III A aluminum…3 valence electrons
IV A carbon…4 valence electrons
V A nitrogen…5 valence electrons
VI A oxygen…6 valence electrons
VII A fluorine…7 valence electrons
VIII A neon…8 valence electrons