The following lesson was originally from the website electronconfiguration.info, by Lynda Jones. That website no longer exists. It has been excerpted and adapted for use here,
What is an Electron Configuration?
An electron configuration is a shorthand description of how electrons are arranged around the nucleus of an atom.
They help us predict chemical behavior. We can predict whether two elements will react or not, and if they react, we can also predict what kind of reaction we are likely to have, as well as how strong the reaction will be.
The Anatomy of an Atom
(1) The nucleus in the center of the atom which takes up almost no space at all, and
(2) the electron cloud which occupies all the remaining space in the atom.
The nucleus is so tiny compared to the actual size of the atom that if we were to draw it to scale on this page, you could not even SEE the nucleus.
The Vastness of the Electron Cloud
Now that you have a picture in your mind of just how vast an atom’s electron cloud is, let’s learn how to write electron configurations.
Not Orbits, but Orbitals
We used to imagine that electrons were tiny solid objects, orbiting around the nucleus just like a plant revolves around the Sun. We imagined it worked like this:
See why we called the path of the electrons “orbits”? Welp, although electrons are real, they’re not solid, nor do they orbit the nucleus. If you really want to know what electrons are and how they behave, you can read our articles on early development of quantum theory and quantum mechanics . But for now we’ll think of electrons as clouds of energy that fill up a space.
We still loosely use words like “orbit” or “orbital”, but we don’t take them literally. Nowdays, when we talk about “orbitals”, we really just mean “the space that an electron cloud takes up.” That shape depends on its energy, orientation, and angular momentum.
These electron orbitals take on various shapes – and they can be quite beautiful. Here are the basic shapes that exist.
Precisely describing an electron’s orbital requires calculus, differential equations, and college level modern physics. But – wait for it – most of us don’t need all that.
Turns out, that if we learn a simple naming pattern, we can describe the “orbitals” of electrons around any atom. That let’s folks figure out which atoms will or won’t react with each other.
Pattern in the Periodic Table
You probably already know that the elements of the periodic table are organized by number from 1 – 117. The number of the element, called the atomic number, is the number of protons found in the nucleus of exactly one atom of that element. The atomic number is also the number of electrons an atom has WHEN THE ATOM IS NEUTRAL.
Some of the outer electrons in an atom are able to jump off of one atom, and onto another. So the number of electrons an atom has can change.
The protons, however, are found deep inside the atom, in the nucleus, and CANNOT come and go. Therefore the # of protons fixes the identity of the element, regardless of the number of electrons.
Let’s look now at the Periodic Table.
As we study electron configuration (the arrangement of electrons in atoms), we discover a pattern. Study the periodic table below and see what you can discover.
Notice the s, p, d and f blocks. Notice the sequence of main energy levels. Notice the superscripts which tell how many electrons are in each orbital sublevel.
The pattern becomes more complex with each succeeding layer, and we have not included all the details here, but you should still be able to recognize a general trend. We have chosen to place in each element box the one term in the electron configuration pattern which shows movement in that section. In some cases, this is the last and highest term in the configuration. In other cases it is not.
The Standard Electron Configuration Pattern
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10
Is there a pattern here?
The large numbers in the sequence, i.e., “1, 22334, 345, 456, 4567, 56,” are called “coefficients.” These numbers refer to the MAIN energy levels of electrons in the atom and are determined by the electrons’ average distance from the nucleus.
The letters, i.e., “s, spsps, dps, dps, fdps, fd,” refer to the orbital sublevels in which the electrons are found. Each MAIN energy level has a number of sublevels equal to its value. If you look more closely at the sequence again,
you will notice the following pattern:
MAIN energy level “1” has only 1 sublevel, the “s” orbital sublevel.
MAIN energy level “2” has two sublevels, both the “s” and “p” sublevels.
MAIN energy level “3” has three sublevels, “s”, “p” and “d.”
The pattern continues with MAIN energy levels “4, 5, 6 and 7,” but the higher sublevels of “5, 6 and 7” (those shown in parentheses above) do not appear in the standard electron configuration pattern, because there are only 117 known elements, so after 117, we don’t have any more electrons to put into the higher orbitals.
Each succeeding MAIN energy level has more orbital sublevels because as we move out from the nucleus of the atom, there is more room for more electrons which can have the same MAIN amount of energy but have different orientations in space and different angular momenta.
Now look at the small black raised numbers.
The small raised numbers shown above in black are called “superscripts.” Superscripts represent the maximum numbers of electrons each orbital sublevel can hold. An orbital sublevel may hold LESS electrons than its maximum, but it can never hold MORE.
How to write Electron Configurations
Once you can write the standard electron configuration pattern, you can then write the electron configuration for any atom or ion. (An ion is an atom with an unequal number of protons and electrons.) To write the electron configuration, you just count the number of electrons you have and use as many orbital sublevels as you need to hold all your electrons. Be sure to always fill the lowest energy levels first.Once you can write the standard electron configuration pattern, you can then write the electron configuration for any atom or ion. (An ion is an atom with an unequal number of protons and electrons.) To write the electron configuration, you just count the number of electrons you have and use as many orbital sublevels as you need to hold all your electrons. Be sure to always fill the lowest energy levels first.
For example, let’s say you have an atom of lithium.
Lithium’s atomic number is 3. So a neutral atom of lithium has 3 protons and 3 electrons. We would need space for 3 electrons. Write:
Now let’s try flourine
What if we had an atom of arsenic, with atomic number 33? We would fill up all the orbital sublevels in the electron configuration pattern until we got to the 4p orbital sublevel. Then we would have only 3 electrons left. 2 + 2 + 6 + 2 + 6 + 2 + 10 + 3 = 33. The electron configuration for arsenic, then, is as follows:
Now, what if lithium were an ion, rather than an atom? (Remember, ions have unequal numbers of protons and electrons.)
The lithium ion has a +1 charge. The lithium ATOM has 3 electrons. How many electrons does the lithium ION have? Remember, electrons come and go, because they are on the “outside” of the atom, but the protons, which are held tightly together in the nucleus of the atom just stay there. So with 3 protons, how many electrons would we need to have in order to get a +1 charge? We would need a number of electrons which is 1 less than the number of protons. Answer: 2.
3 positive charges (from the protons) and 2 negative charges (from the electrons) add up to a total of +1. So we have only 2 electrons, and the electron configuration for the lithium ION is:
How about a fluorine ion? As an ion, does it gain electrons or lose electrons? Its closest noble gas is neon, with 10 protons and 10 electrons. The elements always tend to be most stable in the “noble gas electron configuration.” So fluorine would take on one EXTRA electron to have a total of 10 electrons like neon. 2 + 2 + 6 = 10.
9 positive charges (from the protons) and 10 negative charges (from the electrons) add up to a total of -1. So we have 10 electrons and the electron configuration for the fluorine ION is:
Comparing Arsenic atom with Arsenic ions
Because of arsenic’s position on the periodic table, it can make four different ions:
As+3, As-3, As+5 and As+1
Electron configurations for the first two ions of arsenic may be written using the basic rules we have already learned.
Overlap of Orbital Sublevels
In order to understand and predict the electron configurations for ions such as As+5 and As+1, we need one more piece of information about how electrons arrange themselves in atoms. Take a closer look at the standard electron configuration pattern. For your convenience, the first part of the standard pattern is reproduced here.
Do you notice that the progression of the third MAIN energy level (in yellow) is “interrupted” by the fourth MAIN energy level (in green)?
As you get further away from the nucleus, the sublevels begin to overlap each other. Of course, there is a pattern to the overlap, just as there is a pattern for everything else, but where this overlap begins, writing electrons configurations requires a bit more knowledge about how electrons in atoms behave.
The Rule of Symmetry
Atoms or ions with complete or symmetrical arrangements of electrons are more stable than atoms or ions with incomplete or asymmetrical arrangements of electrons.
Think of a washing machine. When a washing machine’s load is balanced during the spin cycle, everything runs quietly and smoothly, but if the load is not well balanced, the machine begins to thump loudly and rock wildly, acting almost as if the machine were going to break. It is the same way with electrons arranged around the nucleus of an atom. When an atom’s electrons are symmetrically arranged, the atom is in a low energy state and fairly stable, but if the electrons are NOT symmetrically arranged, the atom is unstable and will tend to do one of three things: (1) It may react to form a bond with another atom. (2) It may give up some of its wild outer electrons to obtain a more symmetrical distribution of electrons, even though it will end up as an ion with a positive charge. Or, (3) It may take on extra electrons to gain a more symmetrical distribution of electrons, even though it will end up as an ion with a negative charge.
It is the drive for symmetry, stability and lower energy which drives basically all chemical reactions.
Applying the Rule of Symmetry to Writing the Electron Configuration for the Arsenic +5 Ion
We will now study the Rule of Symmetry by looking at the next two ions of arsenic.
Do you notice in the electron configuration for the As +5 ion shown above that both the 4s2 and 4p3 terms (which are in green in the electron configuration of the neutral atom) are missing?
As the 3rd energy level (yellow) gets completely full, it becomes spherically symmetrical. All electrons from the 3rd energy level are evenly distributed around the atom from the perspective of the atom’s nucleus. Because of its completeness and symmetry, the 3rd energy level resists being broken down when electrons are removed from the atom, so electrons are taken from the higher energy 4th level..
Even though according to the pattern in the periodic table the 4s sublevel gets filled before the 3d sublevel, when electrons are removed from atoms, they are taken first from the highest available MAIN energy level, which, in the case of arsenic, is the 4th level.
That is because electrons in the highest MAIN energy level have the highest energy overall, are on the average at a greater distance from the nucleus, and destabilize the atom the most by whizzing around at higher speeds in an unsymmetrical way. You may imagine these high energy electrons “tugging” on the atom in a manner similar to the way an unbalanced load “tugs” on the center post of a washing machine during the spin cycle.
The As+5 ion with the electron configuration shown above is more symmetrical, and hence more stable, with its completely filled 3rd MAIN energy level than it would be with both the 3rd and 4th level each only partially filled. This is true, even though it is missing five “4th-level” electrons and ends up as an ion with a charge of +5!
There are 3 names for the l quantum number: azimuthal, angular, and orbital.
…The different orientations are indicated with the m quantum number. This is also called the magnetic quantum number. Notice the m quantum number starts with the negative of the l quantum number. Since we are at the l=1 quantum number (p orbital), the m quantum numbers start out as -1, then go up to 0 on the next electron, and +1 on the third electron…
Practice Problem 8: How many electrons are in this element that has this electron configuration? 1s22s22p3