Electron Shells

Picture an atom. Of course, you’ll probably never have seen one properly before - they’re tiny, tiny things. Take the thickness of a standard piece of printer paper, for example. How many atoms thick do you think it is? One thousand? Fifty thousand? Two hundred thousand? The answer is one million. Yes, really - one million atoms are only just as thick as a sheet of paper. In fact, it would take one hundred million atoms to form a line just one cm long.

An atom, as you’ll remember from Fundamental Particles, contains a nucleus full of protons and neutrons. This nucleus is extremely small and extremely heavy. If our atom was the size of a football stadium, the nucleus would only be the size of a marble. Most of what is left in the atom is empty space, but our atom also contains electrons, orbiting around the nucleus in things known as shells. These electron shells are an important part of electron configuration and atomic structure, and play a role in determining an atom or ion’s reactivity. But what exactly are they?

Each electron shell is given a number according to its distance from the nucleus, called the principal quantum number, n. Principal quantum numbers start at 1 and increase by 1 each time, so the first four energy levels have the principal quantum numbers 1, 2, 3, and 4 respectively. The higher the principal quantum number, the higher the energy level of the shell and the further away it is from the nucleus.

Higher energy shells can also hold more electrons. The first shell can only hold two electrons, but the second eight and the third eighteen. The general rule for the number of electrons a shell can hold is , where n is the shell’s principal quantum number. For example, the second shell can hold electrons.

Fig. 1 - A diagram showing how quantum numbers link to distance from the nucleus. As their quantum number increases, electron shells get further from the nucleus and can hold more electrons

What are electron sub-shells?

Electron shells are split into smaller sub-shells which themselves contain orbitals. We’ll explore subshells together first before moving on to orbitals.

Sub-shell types

Each energy level, which you’ll remember is just another term for an electron shell, contains a certain number of sub-levels. These are also known as sub-shells. You can think of sub-shells as mini divisions within each shell or energy level. The first four types of sub-shell are s, p, d, and f.

However, not every shell contains each type of sub-shell. For example, the shell nearest the nucleus with n = 1 only contains an s sub-shell. We call this sub-shell 1s. The second shell contains sub-shells 2s and 2p, while the third shell contains 3s, 3p, and also 3d.

Sub-shell energy levels

We know that each electron shell has its own energy level. As the principal quantum number increases, the shell increases in energy level. Likewise, each of the sub-shells within a shell has a different energy level too. S sub-shells have the lowest energy level, then p, then d, then f. But you should remember that all the sub-shells in one electron shell have a lower energy level than the sub-shells in an electron shell with a higher principal quantum number. That may sound a little confusing, but it simply means that all the sub-shells in energy shell 2, for example, have a lower energy level than the ones in shell 3. However, there is just one exception. Sub-shell 3d has a higher energy level than 4s, despite it being in a shell with a lower principal quantum number.

Fig. 2 - Electron Shells quantum number energy shells StudySmarter

What are electron orbitals?

Let’s look at hydrogen. You’ll remember that hydrogen has one electron (see Atomic Structure), and if you plot this electron’s location again and again, you’ll eventually end up with a sketch looking something like this:

Fig. 3 - Hydrogen's one electron mostly occupies a spherical region

We know this region as the orbital from the sub-shell 1s. As you can see, this orbital is roughly spherical. Let’s take a closer look at the shapes and properties of all the other orbitals.

Orbital shapes

Orbitals have different shapes, depending on their sub-shell. S orbitals are spherical, p orbitals are a figure of eight, and d orbitals can have a variety of shapes.

Fig. 4 - A diagram showing the shapes of the s orbital, left, and the p orbital, right

Electron spin

The electrons in an orbital must have opposite spins. Spin is a property of electrons that can take it either up or down. In an orbital there can be at most one electron with an up spin, and one with a down spin. (Explore spin more in Understanding NMR.)

Orbital energy

Orbitals in the same sub-shell all have the same energy. That means, for example, that all 10 electrons in the 3d sub-shell have the same energy as each other; both of the electrons in 2s have the same energy as each other.

The following diagram puts together what we know about shells, sub-shells, orbitals, and energy levels to show the quantities and energies of the orbitals up to 4p.

Fig. 6 - A diagram showing the energies of the different electron shells, sub-shells, and orbitals. Remember that each orbital can hold up to two electrons

Electron configuration

Electrons fill shells, sub-shells, and orbitals in a certain order. They’re quite fussy, really - they like following certain rules. Take a look at Electron Configuration to find out more about how exactly electrons are arranged in an atom, but for now you should know that an atom’s electron configuration determines its reactivity and properties.