Energy-Level Diagrams and Electron Configurations
Its good to know we CAN determine quantum addresses for each
electron and that we CAN draw a pic of these atoms (handout). But,
most of what we need to explain with this model can be done with EL
diagrams.
 Rules for drawing Energy Level diagrams:
1. Pauli exclusion principle: no 2 electrons can have the same 4
quantum numbers
2. aufbau principle: electrons are added to
the lowest energy orbit first
3. Hund’s rule: electrons occupy each of several orbits in a subshell
before pairing up. Half filled orbitals and full orbitals are
specially stable.
Example:
Electron configuration
Eg. oxygen - O: 1s22s22p4
Eg. Iron
Fe: 1s22s22p63s23p64s23d6
Eg. Helium
Condensed electron configuration
Eg. Oxygen O: [He] 2s22p4
Eg. Iron
Fe: [Ar] 4s23d6
Eg. Antimony
What can we explain with this Model?
1.
Trends in Ionization Energy and Atomic Radius
Read the two page handout and answer these guiding questions….
 Does our quantum model explain the trend in AR? How?
 Can the model explain why there are some ‘anomalies’ in the
transition metals radii?
 What is the general trend going across a period or going down a
group in IE? Why?
 Can the quantum model explain why phosphorus is unusually
high IE? Or why Sulfur is unusually low?
2.
Anomalous Configurations
 Predict the electron configuration for Copper and for Silver
Write them here:
 Now look at the ACTUAL configurations…
Remember – half filled and full filled orbitals are specially stable…
3.
Ions of the Transition Metals
 For negative ions (like S2-), draw the energy level diagram and then
add the required (example 2) electrons to the levels in order
Try to predict the common charge on the halogens
For positive – draw the neutral atom’s energy level diagram, then take
away from the highest quantum number first (so 4s before 3d); also, ½
filled orbits are special
Try to predict WHY iron can be +2 or + 3
4.
Paramagnetism
 ½ filled orbits are associated with being paramagnetic
Why is iron paramagnetic?