Electron Configuration

Electrons play a crucial role in
chemical reactions
and how compounds interact with each other. Remember,
electrons are the negative particles in an atom that “orbit” the nucleus. Although
we say they orbit the nucleus, we now know that they are actually in a random state
of motion surrounding the nucleus rather than making circles around it, which is
what an orbit implies. The best analogy to describe electron motion within an atom
is how bees buzz around a beehive. They don’t fly in complete circles around it,
but they do hover and move around it in a seemingly random motion.

Electrons increase in elements as protons do, which is from left to right and from
top to bottom on the periodic table. Therefore, the element with the fewest electrons
would be in the top left-hand corner of the table and the element with the most
electrons would be in the bottom right hand corner. The elements are arranged so
that the increase from element to element is one electron. Therefore, in the first
row, we see hydrogen and helium. This is because hydrogen has one electron and helium
has two electrons, so we place them in ascending order.

Electron Orbitals

We categorize electrons according to what orbital level in which they reside. The
four orbitals are s, p, d, and f. They are classified by divisions on the periodic
table, as follows:

The first orbital is the s orbital. It has room to hold two electrons. The electrons
have opposite spins, so it makes sense that they are paired together. The s orbital
is a sphere, with the x, y, and z axes passing through it, like this:

This means that the two electrons can occupy any of the space seen in this sphere,
and they sort of “hover” around in the given space.

The next orbital is the p orbital. It can hold up to six electrons, therefore it
has three sub-orbitals (each can hold two electrons). The spins on electrons are
still opposite, this time split into three and three (since the first orbital only
held two electrons, we said the spins were opposite. Now that this orbital can hold
six electrons, three spin one way and three spin the opposite way). The p orbital
is not sphere shaped, however it does have six lobes that are shaped like balloons.
Two lobes are on the x axis, two are on the y axis, and two are on the z axis. These
three separations are considered sub-orbitals and combine to make up the entire
p orbital. The nucleus of the atom is located where these three axes meet. The p
orbital looks like this:

The next orbital is the d orbital. It can hold up to 10 electrons, therefore it
has five sub-orbitals (each can hold two electrons). The spins of the electrons
are opposite, so five are spinning one way and the other five are spinning the opposite
way. The d orbital is not sphere shaped; it looks more like the p orbital, except
there are more lobes that cannot be shown all at once. We showed the entire p orbital
(all three of the sub-orbitals) in one diagram, because there were two lobes on
each axis. However, we need to show the five different sub-orbitals of the d orbital
in order to fully explain where the lobes are located, and how they are shaped.
We will show you four views, with labels on all of the axes.

The first view is of the lobes that lie on the XY plane, shown in aqua here. The
second view is a three dimensional view of lobes on the Z axis that rotate 360 degrees
around the axis. There are two lobes, one in the top hemisphere and one in the bottom,
and a tube-shaped area that circles the Z axis and intersects the X and Y axes.
It’s shown here in orange. The third view is of the lobes on the ZY plane, with
the X axis running perpendicular to it. It’s shown here in green. The last view
is of the lobes lying on the ZX plane, and is shown here in pink. If all of these
layers were put together, we would see a sort of star-burst image, with a tube encircling
the middle.

The final orbital is the f orbital, and scientists are not completely sure of the
shape of its orbital. However, they do have seemingly accurate predictions of where
electrons will fall. We will show you the following probabilities of where electrons
lie:

We showed you two probabilities of where the f orbitals lie; however, the first
image (in blue) is shown on the Z axis. It is actually repeated on the X axis and
again on the Y axis. The second image (in orange) is shown in the XYZ dimensions;
however, it is repeated three more times for a total of four positions using this
shape and lobe configuration. We say that these are probable locations because scientists
cannot actually track and determine the exact location of electrons. However, through
research and abilities to track electrons in other orbitals, scientists can say
that the likely location of f-level electrons is in one of these locations.

Diagonal Rule, or Madelung’s Rule

In chemistry, the Diagonal Rule (also known as Madelung’s Rule) is a guideline explaining
the order in which electrons fill the orbital levels. The 1s2 orbital
is always filled first, and it can contain 2 electrons. Then the 2s2
level is filled, which can also hold 2 electrons. After that, electrons begin to
fill the 2p6 orbital, and so on. The diagonal rule provides a rule stating
the exact order in which these orbitals are filled, and looks like this:

As you can see, the red arrows indicate the filling of orbital levels. Starting
at the top, the first red arrow crosses the 1s2 orbital. If you follow
these arrows down the list, you can easily determine the order that electrons fill
the orbital levels.

There is an exception to this rule when filling the orbitals of heavier electrons.
For example, when filling the 5s2 orbitals, the rule says that 5s2
will fill, and then 4d10 will fill. However, when filling these orbitals
for certain metals, only one electron will fill the 5s2 orbital, and
the next electron will jump into the 4d10 orbital. This can be predicted,
but cannot be exactly determined until it is observed. The same is true for the
6s2 orbital-for certain heavy metals, the 6s2 will only contain
one electron, and the other electrons will jump to the 5d10 orbital.

Electron Notation

Following the diagonal rule, there is an easy way to write electron configuration.
We are simply going to use the orbital names we learned from the diagonal rule (1s2,
2s2 and so on). However, we are only going to write the number of electrons
that the atom actually contains. For example, hydrogen has one electron, which would
fall in the 1s orbital. Thus, the electron configuration for hydrogen is 1s1.
We write the superscript as 1 because there is one electron. Helium, the next element,
contains two electrons. They both fill the 1s orbital, so the electron configuration
for helium is 1s2. Here again, we write the superscript as 2 because
there are two electrons.

Electron configuration moves across and down the periodic table. You might have
noticed that we first put one electron in the 1s orbital (with hydrogen), and then
we put two electrons in the 1s orbital (with helium). Continuing this trend, we
would next have 3 electrons with lithium. We would place two of them in the 1s orbital,
and one of them in the 2s orbital, so the electron configuration would be 1s2
2s1. However, we can also write this using the configuration of helium,
because it is a noble gas. Noble gases are stable elements, so we can use their
configurations in determining other configurations. So, instead of writing 1s2
2s1, we would write [He] 2s1. This means that lithium contains
the same configuration as helium, and then has one more electron in the 2s orbital.
Notice that we use brackets to encase the previous noble gas, and then we continue
writing the configuration as we normally would. This might not seem like a big deal,
or a shortcut right now, but once you get pretty far down the periodic table, this
will save you a lot of time and energy.

We’ll show you one short example of this. Let’s say we need to determine the electron
configuration for Ba, Barium. Counting across the table, we would come up with the
following configuration:

1s2 2s2 2p6 3s2 3p6 4s2
3d 10 4p6 5s2 4d10 5p6 6s2

Instead of writing all of this out, we could simply find the previous noble gas,
which is Xe. Therefore, we can write [Xe] and then figure out the rest of the configuration.
We look and see that Ba is in the 6th row, so we know that we’re going to start
with 6s2. We can look and see that Ba is the second element in that row,
so it has two electrons to go in the 6s orbital.

Thus, we can conclude that the final electron configuration for Ba is [Xe] 6s2.

Electron Spin

Every electron placed in an orbital has a feature we refer to as “spin.” We’ve already
talked about electrons having been thought to have a specific “orbit,” and then
later discovered to hover in the places listed above. Well, electrons don’t literally
spin, but their movement sort of looks like someone somersaulting, in a very fast,
random state. This is called spin. Each electron can either have a +1/2 spin or
a -1/2 spin, which indicates its direction of motion. There is never a 0 spin. When
filling orbitals, electrons spin in pairs, one + and one -. The electrons with an
up (+) spin fill first, and the electrons with a down (-) spin fill second. It would
look like this:

and so on. Because the s orbital can hold two electrons, we draw one box in which
two electrons (represented here with arrows) can fit. Since the p orbital can hold
6 electrons, we draw 3 boxes that will each hold 2 electrons. This will continue
with the d orbital (10 electrons fit in 5 boxes) and the f orbital (14 electrons
fit in 7 boxes).

These boxes fill in a certain order-all of the boxes in one column will fill with
up spin electrons first, and then down spin. So, for example, if the element has
a configuration of 1s2 2s2 2p3, it would look like
this:

As you can see, we filled the boxes with up arrows (electrons) first. If we had
more electrons, we would go back and add them in to the second column, as down spin
electrons.

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