Lesson Video: Electronic Configurations Chemistry • 10th Grade

In this video, we will learn about electronic configuration notation, and describe the electronic configurations of atoms and ions for main group elements up to the end of the 4th period of the periodic table.


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.

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