Video Transcript
In this video, we will learn about
electronic configuration notation and describe the electronic configurations of
atoms and ions from main-group elements up to the end of the fourth period of the
periodic table. Electronic configuration notation
is a way of representing the nature of electrons around an atom or ion. It’s an alternative to using the
four quantum numbers. We’re going to recap quantum
numbers and then go into electronic configurations.
The principal quantum number 𝑛
indicates the electron shell. 𝑙, the subsidiary quantum number
or orbital angular momentum quantum number, indicates the subshell. Zero corresponds to s-type
subshells, one to p-type subshells, two to d-type subshells, and three to f-type
subshells. As we look at atoms with more and
more electrons, the range of allowable values for each quantum number increases to
provide more possible unique combinations of quantum numbers. In other words, more addresses
signify more positions in the electron cloud. But we need two more quantum
numbers to get down to the level of individual electrons.
The next quantum number is the
magnetic quantum number, 𝑚 𝑙. The number of allowable values for
the magnetic quantum number for a given value of the subsidiary quantum number
indicates the number of orbitals per subshell: one orbital for an s-type subshell,
three orbitals for a p-type subshell, and five orbitals for a d-type subshell. And the pattern continues with
f-type subshells having seven orbitals.
The final quantum number 𝑚 s, the
spin quantum number, indicates the spin state of an electron, either spin up or spin
down. Plus a half indicates a spin-up
state, and negative a half indicates a spin-down state. Just in case you need them, here
are the general formulas for 𝑙 and 𝑚 𝑙. The key thing to take away is that,
per orbital, there are two possible spin states, so two electrons maximum, either
one, three, five, or seven orbitals per s-, p-, d-, or f-type subshell; 𝑛 minus one
subshells per shell; and the principal quantum number corresponds to the shell
number.
Quantum numbers are very specific,
but they aren’t very condensed. That’s why we use electronic
configurations. This is an electron shell diagram,
and it’s very effective at showing us where electrons are in terms of the
shells. On the other extreme, quantum
numbers are great for telling us the precise state of an individual electron.
To represent this diagram with
quantum numbers, we’d need 28 of them. We can imagine needing many more
for much bigger atoms. Electron configurations provide an
efficient, happy medium. Let’s start building up the
electronic configuration for a nitrogen atom. We start off with the inner shell,
indicated by the number one. There’s only one type of subshell,
an s subshell, in the first electron shell. So we begin our electronic
configuration with the subshell notation 1s. An s-type subshell contains only
one orbital. And an orbital can contain a
maximum of two electrons. In the diagram, there are only two
electrons in the first electron shell. We can place these two electrons in
the 1s subshell, completely filling the subshell and the shell. We can indicate the number of
electrons in a subshell by putting a superscript number next to the subshell
notation.
Now, you might be wondering, where
do we indicate the electron spin? Well, we need a little bit more
information to understand the significance of this notation. We know there are two electrons in
the 1s orbital. So one must be spin up and one must
be spin down. Even though we don’t specifically
account for spin in an electronic configuration, we can figure it out from our other
understanding.
By convention, the first electron
in an orbital is spin up and the second electron is spin down. That’s the first two electrons
accounted for. Let’s move on to the other
five. These electrons are all in the
second electron shell, so we start with a two. In the second electron shell, there
are two types of subshell, an s and a p. By convention, we draw independent
subshell notation for each subshell. For reasons that will become more
important down the line, we don’t write 2sp. We have separate twos for the two
subshells.
The first two electrons go in the
2s subshell. The next three electrons can all go
in the 2p subshell. With three orbitals, there’s more
than enough space. That accounts for all of the
electrons, but we’re not quite sure about the spins yet. The 2s orbital resembles the 1s
orbital, with one electron spin up and one spin down. But in the 2p subshell, it’s not
clear because we have six places and three electrons. They might be like this, like this,
or even like this.
The first principle we need to obey
is that electrons will only start to pair up after all orbitals have at least one
electron. And another principle is that
electrons in the same subshell will arrange themselves so as to maximize their
overall spin. So the first electron is spin up in
the first orbital, the second electron is also spin up in the second orbital, and
the third electron is also spin up in the third orbital. If there were a fourth electron,
that would go spin down into the first orbital.
Each of these principles has a
special name. The first rule is that the
lowest-energy configuration is the one with the most unpaired electrons. In this particular configuration,
there are three unpaired electrons. And we would only reduce that
number if we started to pair them up. The second rule is that these
electrons will have their spins in the same direction, all spin up by
convention. All the unpaired electrons in this
configuration are spin up. These rules, courtesy of Friedrich
Hund, are known as Hund’s rules.
By the way, you’ll often see
electron configurations or electronic configurations. These can be used
interchangeably. It’s more common to see electronic
configurations when we’re talking about the atom or ion and electron configurations
when we’re looking at the notation.
The next thing we’ll look at is how
we know in which order to fill the subshells. The first rule is that the higher
the value of 𝑛 for an electron, the higher its energy. So an electron in a 1s subshell is
more stable than one in a 2s or a 3s subshell for the same atom or ion. The second rule is that, for a
given value of 𝑛, an electron in the s-type subshell is more stable than one in a
p-, d-, or f-type subshell. So an electron in the 4s subshell
has a lower energy than an electron in the 4p subshell.
But one important complication is
that as 𝑛 increases, the range of orbital energies also increases. The energies of the orbitals in the
first and second electron shells are nicely separated. And it’s the same for the second
and third electron shells. But the energy range for the fourth
electron shell overlaps that from the third. This means that some electrons in
the fourth electron shell will actually be more stable than some electrons from the
third electron shell. The equations necessary to predict
the exact energies are incredibly complicated. Thankfully, there’s an easy diagram
we can use to remember the order.
The lowest-energy orbitals occur in
the 1s subshell, followed by those in the 2s and 2p subshells, followed by the 3s
and the 3p. But the 4s subshell is generally
lower in energy than the 3d subshell, although they are very close. So the 4s subshell is filled first
and then the 3d. We could just remember the
sequence, but this particular arrangement helps.
We can use these diagonal arrows to
put the subshells in order of energy: 1s, 2s, 2p, 3s, 3p, 4s, and then 3d. And we can extend this pattern
further to 4p and 5s, 4d, 5p, and 6s, and so on. This tool is very helpful for
figuring out the electronic configurations of atoms or ions. However, there are circumstances
where it doesn’t reproduce reality. In this video, we won’t be taking
things to an extreme where there are significant inaccuracies. But one thing we do need to address
is something we’ve already taken for granted. Electrons will go into the
lowest-energy orbital first. This is known as the Aufbau
principle.
Electrons in a ground-state atom or
ion will occupy the lowest-energy orbital available. You’ll see electronic
configurations in order of energy, which is helpful for knowing exactly where to put
the electrons. Or they may be arranged by lowest
quantum number, for instance, with all the 𝑛 equals three subshells before the 𝑛
equals four subshells. Both presentations are valid,
although you’ll more commonly see the by-energy system when we’re trying to work out
where electrons should go.
Let’s have some practice of
populating subshells using krypton as an example.
The element krypton has an atomic
number of 36. Therefore, we know the krypton atom
has 36 protons in its nucleus and 36 electrons in its electron cloud. The first two electrons will go
into the 1s subshell, that’s the first two electrons accounted for, two more coming
to the 2s subshell. The next subshell to fill is the 2p
subshell. p-type subshells contain a maximum of six electrons. We have more than enough to fill
this subshell. So the next six electrons go
here.
Next, we have the start of the
third electron shell, with the 3s and 3p subshells. That’s two electrons in the 3s
subshell and six in the 3p. We don’t populate the 3d subshell
until we fill the 4s subshell. So that’s two more electrons down,
followed by another 10 in the d-type subshell. d-type subshells have five orbitals
and have a maximum occupancy of 10 electrons. To figure out where the remaining
six electrons go, we need to continue our diagram.
Next stop is the 4p subshell, which
can contain a maximum of six electrons. That’s exactly the number of
electrons that we have left. With all the electrons now
accounted for, we know that the electronic configuration of a krypton atom in its
ground state is 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6.
As you can see, electronic
configurations can get quite long. We can simplify electronic
configurations by using brackets placed around element symbols. For instance, brackets around the
element symbol for krypton indicate this full electron configuration. For instance, we can represent the
electronic configuration of a rubidium atom, which has one more electron than a
krypton atom, in this way. This way, we can focus on the 5s
electron, which is more important for understanding the chemistry of rubidium than
the core electrons. Typically, we only see group 18
elements in this notation, helium, neon, and argon being good examples.
Next, let’s have a look at
ions. When figuring out the electronic
configuration of ions, we look first at the electronic configuration of the
equivalent atom and add or take away electrons. We remove electrons from the
highest-energy occupied orbital or add electrons to the lowest-energy available
orbital. The electronic configuration for a
sodium atom is 1s2 2s2 2p6 3s1. In this configuration, the 3s
subshell has the highest energy. Therefore, when we remove an
electron from a sodium atom, the electron in the 3s orbital will be removed
first. The electronic configuration of a
sodium plus ion can be written as 1s2 2s2 2p6, ignoring the 3s subshell
completely. But it’s perfectly valid to write
3s0 to highlight which subshell the electron has been lost from relative to the
atom.
You may have noticed that 1s2 2s2
2p6 is the same electron configuration we see for an atom of neon. We can express the commonality
between sodium plus and neon by saying they are isoelectronic. What we mean by this is that an
atom of neon has exactly the same number of electrons as an ion of sodium. It’s very important to understand
that electronic configurations can be assigned to atoms or ions. If we only have the electronic
configuration, we can’t tell for sure whether we’re looking at a sodium plus ion or
a neon atom, or any one out of a massive set of anions and cations. To work out which we’re talking
about, we need to be told we’re dealing with the atom or an ion with a specific
charge.
Let’s have some practice reading
electronic configurations. Here’s an example: 1s2 2s2 2p6 3s2
3p4. And we’ve been told it corresponds
to an atom. We can work out the total number of
electrons by adding the occupancies of each subshell. From this, we can determine that
this atom has 16 electrons in total. Atoms are, by definition, neutral,
having the same number of protons as electrons. This particular atom must therefore
have 16 protons. We can determine the element for
this atom by looking at the number 16 on the periodic table. The element with atomic number 16
is sulfur. Therefore, this configuration is
for a sulfur atom.
The next example is a configuration
of 1s2 2s2 2p6 3s2 3p6 for a two plus ion. We start by doing exactly the same
thing, adding up the number of electrons in the electronic configuration. However, to form a two plus ion
from an atom, the atom needs to lose two electrons. Therefore, to work out the number
of electrons in the atom, we have to add them back. Therefore, in the equivalent atom,
there would be 20 electrons, 20 protons, meaning that we must be dealing with an ion
of calcium. The calcium two plus ion has the
electron configuration 1s2 2s2 2p6 3s2 3p6.
When we want to quickly figure out
the electronic configuration of an atom of an element in its ground state, there’s a
trick we can do with the periodic table. The periodic table is essentially
arranged based on the electronic configurations of the atoms of the elements. If we look at the position of the
last electron in the electron configuration of the atoms, we get an interesting
pattern. The outer electron of a hydrogen
atom is in the 1s subshell, and it’s the same for an atom of helium. And we see a similar pattern for
lithium and beryllium. We can extend this for boron to
neon and further.
If we want to quickly recall the
electronic configuration of an atom of neon, we just need to account for all the
preceding electrons. There are the two electrons in the
1s subshell, the two electrons in the 2s subshell, and six electrons in the 2p
subshell. This trick is pretty good for
elements in the s or p blocks, but it’s not as effective for elements in the d or f
blocks.
Let’s finish up with the key
points. Electronic configuration notation
is an effective way of condensing the information about the state of electrons in an
atom or ion. The notation is a sequence of
subshell symbols with superscripts to indicate the number of electrons in that
subshell. The subshells can be ordered by
energy or by quantum number. Whatever way the subshell notation
is arranged, the subshells are filled in order of lowest energy first.
We can make an estimate of which
subshell has the lowest energy by arranging them in an Aufbau diagram, where
subshells are ordered by 𝑛 and 𝑙. This estimate is pretty effective
for the main-group elements. And Hund’s rules tell us that
electrons in the same subshell will only pair up if there are no empty orbitals and
that unpaired electrons in the same subshell will have the same spin.