This entire lesson is from : http://electronconfiguration.info/
How do you
Write Electron Configurations?
(Site updated: November 16, 2010)
Remember, chemistry IS easy! (If you have the right tools to help you learn it.) The “Purney Cheer” created by Mr. Ron Purney, a chemistry teacher in Ohio, and reproduced here with permission, is the BEST tool I know to learn how to write electron configurations. Read on!
What is an Electron Configuration?What is an Electron Configuration?
The Anatomy of an Atom
An atom may be divided into two sections: (1) The tiny 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 spherical space in the atom. The nucleus of the atom 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. And if we kept the nucleus the size you see it in the representation below (approximately 6 mm across), then the TRUE size of the corresponding electron cloud (represented in green) would reach outward in each direction approximately 10 thousand times that distance, or 60 meters (~ 197 feet). If the atom were drawn to scale with the nucleus the size you see it here below, its spherical electron cloud would have a diameter of 120 meters, longer than the length of an entire football field! (A football field is 360 feet long or 110 meters.)
The Vastness of the Electron Cloud
To get a sense of just how vast the electron cloud is compared to the size of the atom’s nucleus, look at the black circle in the middle of the drawing below and imagine that to be the size of your nucleus. (It should be about 6 mm across.) Then imagine the little black nucleus in the middle of a football field. In your mind, rotate the entire football field around the nucleus so that it swings far outside the stadium on either side, way high into the air above it and deep down into the earth beneath it, and until you can visualize a perfect sphere that large. Hopefully this gives you some idea of just how much of an atom is really “mostly empty space.”
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
Since all electrons do NOT orbit the nucleus per se as first believed, scientists now use the word “orbital” to describe the shape of the space an electron occupies in an atom, based on the combination of its energy, orientation in space and angular momentum. Describing exactly where electrons are found in an atom is a mathematical exercise so complex that it required an entire10-week course at UCLA just to study it! But you don’t need to understand any of that in order to write a proper basic electron configuration. You simply need to know two things: (1) the standard electron configuration pattern and (2) how many electrons your particular atom has. [A little later, there will be one more piece of information you will need to know, but first let’s start with the basics.]
The Meaning of the Electron Configuration
The word “configuration” means “the arrangement of the parts of something.” An electron configuration describes “the arrangement of electrons in space around a nucleus.” Electrons generally do not orbit the nucleus as scientists used to think. Rather, they hover, wiggle and jiggle, vibrate back and forth, and move in certain patterns based on three factors: (1) their energies, (2) their orientations in space (in other words, exactly where they fit in space relative to the other electrons in the atom) and (3) their angular momentum. (Angular momentum may be thought of as describing how wildly electrons swing about in a circular path around some center point. Some electrons may not necessarily be orbiting completely around the atom but might possibly just be making little circles out to the side, somewhat like if you were making arm circles out to the sides of your body.)
As you study electron configurations, you will see that there is a wonderful pattern. The pattern of electron configurations of the elements explains perfectly why the periodic table is organized as it is.
The Electron Configuration 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. It is very important to know at this point that some of the outer electrons in an atom are able to jump onto or off of other atoms according to certain physical laws, so the number of electrons an atom has can change from moment to moment. The protons, however, are found deep inside the atom, in the nucleus, and CANNOT come and go. Therefore, the number of protons in an atom fixes the identity of the element, regardless of the number of electrons an atom has from one moment to the next.
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. For an authoritative list of complete configurations, especially for the more complex higher elements, go to the Israel Science and Technology Home Page.
The Standard Electron Configuration Pattern
The standard electron configuration pattern is as follows:
You might think this is difficult, but it’s really not. A high school chemistry teacher friend of mine, Mr. Ron Purney from Ohio, created a series of rhythms which makes it SO EASY to do electron configurations that all his students get A’s in this topic every single time. Here is the secret.
How to Quickly Memorize the Standard Electron Configuration Pattern using:
The Purney Cheer
(1) Look at the pattern of numbers:
1, 22334, 345, 456, 4567…, 56.
If you SAY these numbers, pausing for each comma, with a double pause before the last “56,” you will notice an easy rhythm. Memorize this sequence.
(2) Next, look at the pattern of letters:
s, spsps, dps, dps, fdps…, fd.
Again, say these letters, pausing for each comma, with a double pause before the last “fd.” Memorize this sequence, too.
(3) Now write down the numbers first, leaving space in between, like this:
(4) Then match the letters to the numbers by filling in the spaces leaving a little bit of space after each letter for the superscripts (small raised numbers).
(5) Finally, after each letter write the corresponding superscript.
You have just reproduced the entire electron configuration pattern which can be used for all the elements all the way up through the last element of the periodic table!
To help you memorize this pattern, you may want to download “The Purney Cheer.” It is available as a PowerPoint presentation for teachers and students for only $0.99. [Not yet active.]
Analysis of the Standard Electron Configuration Pattern
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 Use the Standard Electron Configuration Pattern to write Electron Configurations for Different Elements and Ions
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.
There are 2 electrons in the 1s sublevel and 1 electron in the 2s sublevel. 2 + 1 = 3. That’s it!
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.
The 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. So you can get an idea of this spread and overlap of main energy levels and orbital sublevels in an atom, study the graphic below.
There is an extremely cool music video, by the way, which can help you understand and remember the pattern for orbital sublevel overlap, and it is SO EASY to learn, you won’t believe it. You can follow the video using your hands and you simply won’t forget it. To download this video, click here. [Not yet active.]
The Rule of Symmetry
The Rule of Symmetry is this: Atoms or ions with complete or symmetrical arrangements of electrons are more stable than atoms or ions with incomplete or asymmetrical arrangements of electrons. As an analogy, 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.
Here is the electron configuration for the As+5 ion.
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!
To understand the next ion of arsenic, you need to learn this next little bit about the Rule of Symmetry.
The Stability of Completely Filled
Study the following boxes.
Each box contains 4 green egg shapes representing 4 orbitals arranged around a nucleus which is not shown because it is so small, but it is located in the center. The black dots in these drawings represent electrons. Without knowing any more than that, can you tell which of the 8 boxes represents the most symmetrical arrangement of electrons from the perspective of the nucleus in the center?
Hopefully you chose “Box H,” because this is the most symmetrical of all the boxes. Which box would then be the next most symmetrical? Remember, you are thinking from the perspective of a circle, rather than from a line. If you chose “Box D” you are correct.
The more symmetrically arranged electrons are around the nucleus of an atom or ion, the more stable that atom or ion is. A sublevel of orbitals which is completely filled, as shown in “Box H,” is the most stable arrangement. The next most stable is when the sublevel is exactly half filled, as shown in “Box D.” You could also safely say that G is more stable than F, and that F is more stable than E. Similarly, in the top line of boxes, C is more stable than B, and B is more stable than A. But D is more stable than either E or C. In other words, on an energy graph of the energies in different orbital sublevel configurations (shown below), half-filled sublevel “D” is more stable than either “C” or “E.” And completely filled sublevel “H” is the most stable of all. On an energy graph as shown below, the less energy a particular arrangement has the lower on the graph it is drawn and the more stable it is. (You may think of low energy arrangements as having less energy to get away from an atom.)
When Symmetry is More Important than Energy
In losing or gaining electrons, atoms and ions will favor arrangements that have either a completely filled sublevel, such as in Box H or a half-filled sublevel, such as in Box D. Sometimes electrons are “borrowed” from lower or higher energy orbital sublevels to make a preferred sublevel either completely full or half full, because this is more stable.
If we had a choice between Box A and a completely EMPTY orbital sublevel, as shown below, the choice of “completely empty” would be preferable, as it is more centrally symmetrical.
Interestingly, however, only in the case of “s” orbitals, which are already spherical in shape, whether an “s” orbital has 0, 1 or 2 electrons, there is very little difference in its symmetry, because its electrons are free to roam all over the sphere with no restriction.
So when electrons are needed by the atom to fill or half-fill one of its “p” or “d” sublevels, one or both “s” electrons are “borrowed” from the highest “s” orbital available, because whether an “s” orbital has zero, one or two electrons makes very little difference to the atom’s symmetry. This affects electron configurations when there is an overlap of sublevels and an opportunity to borrow, as we will see below.
Borrowing an “S” Electron to Make a “D” Sublevel Half-Filled in the Arsenic +1 Ion
To predict the electron configurations of atoms and ions which do not fall into the simple cases we have discussed previously, we need to know that a completely filled sublevel (Box H) is more symmetrical, and hence more stable than a partially filled sublevel, because the electrons in a completely filled sublevel are evenly distributed around the nucleus. The next most stable electron configuration is when a sublevel is exactly half-filled (Box D). A half-filled sublevel is more symmetrical than a sublevel filled only 30-40% or 70-80%, so an electron configuration with a half-filled sublevel is also slightly more stable.
Applying this to “p” orbitals, a completely filled “p” sublevel, which would have six electrons (p6), is more stable than a “p” sublevel with 4 or 5 electrons. A “p” sublevel with exactly half its maximum number of electrons, in other words three electrons (p3) is more stable than a “p” sublevel with only 1 or 2 electrons. Nevertheless, a completely filled sublevel is more stable than a half-filled sublevel. So p6 would be the most stable, p3 would be the next most stable, and p5 and p4, p2 and p1 would be less stable.
Since the “d” sublevels may hold a maximum of 10 electrons, d10 would be the most stable configuration for that sublevel, and d5 would be the next most stable. Then would come d9, d8, d4 or d3. Keep the rule of symmetry in mind as you attempt to predict electron configurations for atoms with nearly full or nearly half-full “d” sublevels.
In the case of As +1, as shown below, notice that the 3rd MAIN energy level is completely full and that there are a total of 4 electrons in the 4th MAIN energy level. These electrons could distribute themselves as either “4s2 and 4p2” or “4s1 and 4p3.” The first possibility does not make use of the Rule of Symmetry for the 4p orbital sublevel. The second possibility, however, “4s1 and 4p3,” does. The As+1 ion is more stable with the half-filled 4p sublevel than it would be with only 2 electrons in that sublevel, so it “borrows” one electron from the 4s sublevel, which is spherical. This leaves only one electron in the “4s” orbital, making its configuration “4s1.” As stated previously, one electron in an “S” orbital does not destroy the symmetry of the atom or ion. So the electron configuration below is the best and most stable arrangement for the As+1 ion.
This is the end of this discussion on electron configuration. Hopefully this has been helpful. If you see something confusing in this presentation, or if you find errors or have questions, please contact me, Lynda Jones, at sing-smart.com. My goal is to make chemistry EASY! How am I doing?
To Learn Other Chemistry Skills, go to the Chemistry is Easy Homepage.