Lesson 5: Electron Configuration
Now that you have an idea about how the orbitals, sublevels, and main energy levels are ordered, let's consider where the electrons are found. To indicate where the electrons in an atom (or ion) are found, a chemist writes out an electron configuration. You can think of an electron configuration as an adress for the electrons - it shows the energy level, sublevel, and orbital for each electron in the atom. It might seem like this is useless information - but when we look at the periodic table in more depth, electron configurations will help to explain why the table is arranged the way it is and also why elements react the way they do.
Let's take a closer look at Example 19.
All of the orbitals in the same sublevel have the same energy, but different sublevels have different energies. The lowest in the energy is the 1s, then the 2s, the 2p, the 3s, the 3p, the 4s, the 3d, and so on. Ex. 19 shows the sublevels in order of their energies up to the sixth main energy level.
Bohr's model, with the electrons only in main energy levels or "orbits," did not consider that within a main energy level there might be different sublevels with different energies. In his model, all of the orbitals in the same main energy level have the same energy. This is why Bohr’s model worked for hydrogen: it only has one electron, so only the first energy level has an orbital with an electron in it. When more than one electron is present in an atom, the energies of the sublevels change and they are no longer all the same; Bohr's model doesn't work for an atom with more than one electron. The amount by which those energies change depends on the nuclear charge and the number of electrons present. This means that the energies of the orbitals are different for every atom and ion; but it turns out that the order of their energies (which is first, which is second, etc.) is very nearly the same in all atoms and ions.
When electrons "choose" what orbital to be in, they choose the lowest energy orbitals available. To figure out where all of the electrons are in an atom, count the number of electrons, then fill the orbitals, lowest energy sublevels first, until you run out of electrons. This is the "ground state" of the atom.
You have observed atoms in excited states when you looked at the glowing cathode ray tubes and the glowing metal atoms in the bunsen burner flame. One of the odd predictions of wave mechanics is that an electron in a 3s orbital can't jump directly to a 4p orbital, even if the right amount of energy arrives in the form of light. It must jump to a 4s, and from there go to the 4p. This prediction has been verified experimentally.
A list of an atom's sublevels, in order of their energy, showing how many electrons are in each one, is called the atom's electron configuration. We use superscripts to show how many electrons each sublevel contains. A hydrogen atom, for example, has a single electron in the 1s orbital, so its electron configuration would be 1s1.
You will not be required in this lesson to determine the electron configuration of an atom. However, in the next lesson, it will be helpful if you already know what "electron configuration" refers to and have seen some examples.
The first step in doing an electron configuration is to figure out how many electrons the atom (or ion) has. For a neutral atom, this must be the same as the atomic number. For an ion, start with the atomic number, then add a number of electrons equal to the negative charge, or subtract a number of electrons equal to the positive charge. Electrons are then placed into the electron configuration starting with the lowest energy position (the first energy level) and continuing until all the electrons have been used up.
Remember that all s sublevels can hold up to 2 electrons, all p sublevels up to 6 electrons, d sublevels up to 10 electrons, and so on. You must fill a level before you can put electrons into a higher level.
Here is an easy way to determine the energy order of the sublevels. If they are listed as shown, then the energy order can be found by following the diagonal lines from the top down.
This method works more than 90% of the time. As you might guess, the errors involve the higher energy orbitals – the 4s, 3d, and above – since they are closer together in energy.
For our purposes, this method will be sufficiently accurate and you need not learn any of the exceptions. You will not be required to memorize this chart or the correct “filling order” of the sublevels. In the next lesson, we'll learn how to use the periodic table to determine an atom's electron configuration.
Let's look at two more examples before we wrap up this lesson.
Lithium atoms have three electrons. Two can fit in the lowest energy orbital, the 1s. Since each orbital can hold a maximum of two electrons, those two electrons completely fill that sublevel. The third electron must go into the next higher energy orbital, the 2s. Lithium's electron configuration is 1s22s1.
Nitrogen atoms have 7 electrons. The first two fill the 1s orbital, the second two fill the 2s orbital, and the last three go into the 2p sublevel. This sublevel has three orbitals, so can hold up to six electrons. The electron configuration does not say which 2p orbitals the three electrons are in, only that they can be found in the 2p sublevel. (There are two ways to put the three electrons in the orbitals in the 2p sublevel. Two of them can go in one orbital while the third goes in another, or the three electrons can go in three separate orbitals. It shouldn't be too surprising that the electrons are found in three separate orbitals; since electrons repel each other they will each take an orbital rather than share.)
We would say, "The ground state electron configuration of nitrogen: 1s22s22p3." In other words, 2 electrons are located in the 1s sublevel; 2 electrons are located in the 2s sublevel; 3 electrons are located in the 2p sublevel.