Activity 1 - Hund’s First Rule
Question: Can you work out the electron configuration for a) lithium b) chlorine c) zinc
d) vanadium ?
There is one problem with this way of working out the electron configuration: copper and
chromium have one electron in the “wrong” place. This is because electrons gain stability
from being in full or half-full shells. To explain this we need to think about how we actually
put the electrons into the orbitals we identified above. The best way to think about this is to
think of the orbitals as boxes into which we can put the two electrons. Remembering that
electrons repel each other it makes sense that we put one electron in each “box” of a given
energy level before we pair any of them up. This idea is known as Hund’s first rule. An
example of how you would fill up the electrons is shown in Figure 6.
Figure 1. The electronic configuration for iron. Note that all the d orbitals were half filled
before any electrons were paired.
Question: Use this rule to show the arrangement of electrons for a) lithium b) chlorine c) zinc
d) vanadium
Now thinking about copper and chromium, filling the orbitals using the rules we have thus far
would give us 4s23d9 and 4s23d4 configurations respectively. Rather than this being the case
however, one electron from the 4s shell is promoted to the 3d shell to make this, in the case
of chromium, half full or, in the case of copper, full.
Questions: Draw the correct orbital diagrams for copper and chromium