Electron Configurations

In its ground state, the one electron of hydrogen is in principal energy level 1.  The only possible sublevel for this principal energy level is the “s” sublevel which has a spherical shape (just like our firefly example).  We say that hydrogen’s electron is in the “1s” orbital. 

a 1s orbital

If energy were applied to this electron, say from an electric current, it may move up to another place farther away from the nucleus.  Let’s say it moved up to principal energy level 2 in the “s” sublevel.  It is now in the 2s orbital.  When that electron loses energy and drops back to its ground state in the 1s orbital, a bit of energy would be given off and would be seen as a certain wavelength (energy, color) of light.  If the same electron were REALLY energized and moved up to principal energy level 3, in the “s” sublevel (the 3s orbital) when it dropped back down to ground state, more energy would be released than before and we would see a bit of light with a shorter wavelength.  It can be seen that even though hydrogen has only one electron, it can move among many different orbitals.  An orbital, then, is best described not as a place where an electron is but a potential place for an electron.  All electrons in all atoms absorb and emit energy in this way.  You can see that for atoms with more electrons, many more possible energy changes are possible.  Let’s now define the orbitals for other atoms besides hydrogen.  

In the helium atom there are 2 electrons and both can fit in the 1s orbital.  We say that the ground state electron configuration for helium is 1s².   (Obviously, for hydrogen, this electron configuration would be 1s¹).  At this point, the 1s orbital is filled.  We cannot fit more electrons so close together since they repel each other.  We have filled the front row at our concert hall.  So, for the element lithium, which has 3 electrons, we must now move farther away from the nucleus where there is more room.  Since we are moving farther away, we go to the next allowed energy level, principal energy level 2 where n=2.  The first sublevel we will fill (remember there are 2 sublevels for principal energy level 2) is again an “s” sublevel which, like the 1s obital, is also a sphere but a larger sphere since electrons are farther away from the nucleus on average.

a 2s orbital

The cumulative ground state electron configuration for lithium is then 1s²2s¹.  We can always determine what atom we are dealing with in an electron configuration by adding all the subscripts in an electron configuration.  In this case, the subscripts add up to “3” which is how many electrons lithium has.  The next element, beryllium, has 4 electrons and the last one will be in the 2s orbital giving a ground state electron configuration of 1s²2s².   It should be mentioned here that the Pauli exclusion principal  states that an orbital can hold a maximum of 2 electrons.  This rule arises from experimental data that shows that electrons are spinning on their axes, just like the earth is.  This spinning negative charge creates a magnetic field - the electron is like a tiny magnet.  If 2 electrons occupy an orbital, they have opposite spins and their magnetic fields cancel.   An orbital cannot have 2 electrons with the same spin or their magnetic fields would repel, thus a third electron in an orbital is never found. 

So, where does the 5^th electron in the element boron go?   Principal energy level 2 has 2 sublevels and the “s” sublevel is filled, so the “p” sublevel now will accept electrons.  The “p” sublevel is actually made up of 3 separate orbitals, each holding 2 electrons in shapes that looks like dumbbells. Each dumbbell is oriented along  3 different axes (x,y,z) and is denoted p_x, p_y and p_z. 

a 2p orbital	all 3 2p orbitals

These 3 dumbbell-shaped orbitals can hold a total of 6 electrons.  The p_x, p_y and p_z distinctions are usually not used in electron configurations.  Ground state configurations for the next 6 elements are summarized below:

At this point, all of the possible places for electrons in the first 2 principal energy levels are filled.  A total of 10 electrons can fit in these energy levels, 2 in energy level 1 and 8 in energy level 2.  We must move farther from the nucleus to place the next electron.  Sodium has 11 electrons and its 11^th electron will fit into principal energy level 3, in the “s” sublevel.  Hopefully a pattern is emerging – we always begin a new principal energy level by putting electrons in the “s” sublevel.  The “s” sublevel of any given principal energy level has slightly lower energy (and is thus closer to the nucleus) than its corresponding “p” sublevel.  Again, we can fit 2 electrons into the 3s orbital which fills it and we must move to the 3p orbitals where we can fit a total of 6 more electrons.  So, starting with sodium, we can write the ground state electron configurations for the next eight elements:

You may remember that principal energy level 3 contains 3 sublevels or shapes s, p, and d.  The d sublevel actually consists of 5 separate orbitals holding 2 electrons each for a total of 10 electrons.  These shapes have multiple lobes and are oriented both between and along the x,y,z axes with the nucleus at the origin.   As with the p orbitals, we usually don’t differentiate between the electrons in each separate 5 d orbital, but put a total of 10 electrons in them.  Let’s summarize principal energy levels, sublevels and electron capacities to this point:

Note we have added the up to the 4^th sublevel for principal energy level 4 which are called “f” orbitals.  There are 7 separate orbitals in the “f” sublevel holding a total of 14 electrons.  Which elements contain these electrons will be discussed later.       

 Getting back to the electron configurations of the elements, it would make sense that after 18 electrons, the last of which went into the 3p orbitals, the 19^th would go into the 3d sublevel, but that does not happen.   The electron configuration for the next 3 elements is as follows:

What is going on here?  Why do we start to fill principal energy level 4, a higher energy level with electrons supposedly farther away from the nucleus, before we fill the remainder of principal energy level 3 with electrons that are thought to be closer to the nucleus?   To answer this, we must realize that the principal energy levels (n=1,2,3,…etc.) are average energies for all the sublevels contained within that principal energy level.  Each of the sublevels have  slightly different energies with the “s” sublevel being the lowest, and the “p”, “d”, and  “f” sublevels having progressively greater energies.  This causes some overlapping of energies between principal energy levels.   We can see from the following diagram that the 4s orbital has slightly lower energy than the 3d orbitals, so electrons fill there first.  Next, the 3d orbitals are filled and then the 4p orbitals. 

Starting with potassium, the ground state electron configurations for the next 18 elements are as follows:

As we get to elements with greater numbers of electrons, there is quite a bit of overlapping of orbitals between 2 and 3 different principal energy levels.    Without experience, it is difficult to determine the order in which  the orbitals will be filled.  The diagram below, left, shows how this overlapping of orbital energies distorts our nice pattern  we started to see in the electron configurations.  Use the diagram below, right, to determine the order of orbital filling, starting from the bottom.  Each box represents an orbital.  

Based on this order of adding electrons to atoms, which orbital will take the next electron?  It must go into the “5s” orbital, so the ground state electron configuration of the element rubidium is:

The number of valence electrons an element has will become useful when we talk about chemical bonding.

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