A Crash Course in Electron Configurations
As an example to illustrate how the chart works, click on element number 23,
vanadium (V), in the periodic table. It has 23 protons and 23 electrons;
the atomic number increases by one with
each successive element.
I see the 23 electrons in the chart; they're in four different-colored rows,
with three columns labeled s, p, and d.
Good. The higher up an electron is drawn on the chart, the more
energy it has. In fact, the applet can tell you exactly how
much energy you're dealing with: pick any electron and hold the
mouse pointer over it.
If I do that, a little blue number pops up above the electron.
That number tells you, in eV, how
much added energy it would take for that electron to escape from the
nucleus. For the outermost electron in each atom, this value is called the
ionization energy.
The colored rows represent the "main" divisions in energy; they're known as
primary energy levels. You'll notice that vanadium has electrons in
four different primary levels and is located in the fourth row
of the periodic table; this is not a coincidence.
The colored rows aren't exactly straight; there's a little step up between the
s column and the p column, and a bigger one between p and d.
That's right. s, p, and d are called
sublevels; they're smaller "subdivisions" of energy within the
primary levels. You refer to different energy levels using a number for
the primary level plus a letter for the sublevel; for example, you might
speak of an electron in a "3p" state or orbital. Each primary level has one more sublevel than the one
below: the first primary level has only s orbitals, the second has s and
p, the third s, p, and d, and so forth.
If you play around with the applet for a while, you'll discover that each
sublevel has room for a certain number of electrons, and that they don't
always get filled up in the order you might expect. If you'd like to know
more about what's going on, read the heavy
atoms discussion.
