Write Electron Configurations for Atoms of Any Element

An atom's electron configuration is a numeric representation of its electron orbitals. Electron orbitals are differently-shaped regions around an atom's nucleus where electrons are mathematically likely to be located. An electron configuration can quickly and simply tell a reader how many electron orbitals an atom has as well as the number of electrons populating each of its orbitals. Once you understand the basic principles behind electron configuration, you will be able to write your own configurations and tackle those chemistry tests with confidence.

Steps

Assigning Electrons Using a Periodic Table

  1. Find your atom's atomic number. Each atom has a specific number of electrons associated with it. Locate your atom's chemical symbol on the periodic table. The atomic number is a positive integer beginning at 1 (for hydrogen) and increasing by 1 for each subsequent atom. The atom's atomic number is the number of protons of the atom - thus, it is also the number of electrons in an atom with zero charge.
  2. Determine the charge of the atom. Uncharged atoms will have exactly the number of electrons as is represented on the periodic table. However, charged atoms will have a higher or lower number of electrons based on the magnitude of their charge. If you're working with a charged atom, add or subtract electrons accordingly: add one electron for each negative charge and subtract one for each positive charge.
    • For instance, a sodium atom with a +1 charge would have an electron taken away from its basic atomic number of 11. So, the sodium atom would have 10 electrons in total.
  3. Memorize the basic list of orbitals. As an atom gains electrons, they fill different orbitals sets according to a specific order. Each set of orbitals, when full, contains an even number of electrons. The orbital sets are:
    • The s orbital set (any number in the electron configuration followed by an "s") contains a single orbital, and by Pauli's Exclusion Principle, a single orbital can hold a maximum of 2 electrons, so each s orbital set can hold 2 electrons.
    • The p orbital set contains 3 orbitals, and thus can hold a total of 6 electrons.
    • The d orbital set contains 5 orbitals, so it can hold 10 electrons.
    • The f orbital set contains 7 orbitals, so it can hold 14 electrons.
      Remember the order of the letters with this mnemonic:[1]
      Sober Physicists Don't Find Giraffes Hiding In Kitchens.For atoms with even more electrons, the orbitals continue alphabetically past K, skipping letters already used.
  4. Understand electron configuration notation. Electron configurations are written so as to clearly display the number of electrons in the atom as well as the number of electrons in each orbital. Each orbital is written in sequence, with the number of electrons in each orbital written in superscript to the right of the orbital name. The final electron configuration is a single string of orbital names and superscripts.
    • For example, here is a simple electron configuration: 1s2 2s2 2p6. This configuration shows that there are two electrons in the 1s orbital set, two electrons in the 2s orbital set, and six electrons in the 2p orbital set. 2 + 2 + 6 = 10 electrons total. This electron configuration is for an uncharged neon atom (neon's atomic number is 10.)
  5. Memorize the order of the orbitals. Note that orbital sets are numbered by electron shell, but ordered in terms of energy. For instance, a filled 4s2 is lower energy (or less potentially volatile) than a partially-filled or filled 3d10, so the 4s shell is listed first. Once you know the order of orbitals, you can simply fill them according to the number of electrons in the atom. The order for filling orbitals is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s.
    • An electron configuration for an atom with every orbital completely filled would be written: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d107p6
    • Note that the above list, if all the shells were filled, would be the electron configuration for Og (Oganesson), 118, the highest-numbered atom on the periodic table - so this electron configuration contains every currently known electron shell for a neutrally charged atom.
  6. Fill in the orbitals according to the number of electrons in your atom. For instance, if we want to write an electron configuration for an uncharged calcium atom, we'll begin by finding its atomic number on the periodic table. Its atomic number is 20, so we'll write a configuration for an atom with 20 electrons according to the order above.
    • Fill up orbitals according to the order above until you reach twenty total electrons. The 1s orbital gets two electrons, the 2s gets two, the 2p gets six, the 3s gets two, the 3p gets 6, and the 4s gets 2 (2 + 2 + 6 +2 +6 + 2 = 20.) Thus, the electron configuration for calcium is: 1s2 2s2 2p6 3s2 3p6 4s2.
    • Note: Energy level changes as you go up. For example, when you are about to go up to the 4th energy level, it becomes 4s first, then 3d. After the fourth energy level, you'll move onto the 5th where it follows the order once again. This only happens after the 3rd energy level.
  7. Use the periodic table as a visual shortcut. You may have already noticed that the shape of the periodic table corresponds to the order of orbital sets in electron configurations. For example, atoms in the second column from the left always end in "s2", atoms at the far right of the skinny middle portion always end in "d10," etc. Use the periodic table as a visual guide to write configurations - the order that you add electrons to orbitals corresponds to your position in the table. See below:
    • Specifically, the two leftmost columns represent atoms whose electron configurations end in s orbitals, the right block of the table represents atoms whose configurations end in p orbitals, the middle portion, atoms that end in d orbital, and the bottom portion, atoms that end in f orbitals.
    • For example, when writing an electron configuration for Chlorine, think: "This atom is in third row (or "period") of the periodic table. It's also in the fifth column of the periodic table's p orbital block. Thus, its electron configuration will end ...3p5
    • Caution - the d and f orbital regions of the table correspond to energy levels that are different than the period they're located in. For instance, the first row of the d orbital block corresponds to the 3d orbital even though it's in period 4, while the first row of the f orbital corresponds to the 4f orbital even though it's in period 6.
  8. Learn shorthand for writing long electron configurations. The atoms along the right edge of the periodic table are called noble gases. These elements are very chemically stable. To shorten the process of writing a long electron configuration, simply write the chemical symbol of the nearest chemical gas with less electrons than your atom in brackets, then continue with the electron configuration for the following orbital sets. See below:
    • To understand this concept, it's useful to write an example configuration. Let's write a configuration for Zinc (atomic number 30) using noble gas shorthand. Zinc's full electron configuration is: 1s2 2s2 2p6 3s2 3p6 4s2 3d10. However, notice that 1s2 2s2 2p6 3s2 3p6 is the configuration for Argon, a noble gas. Just replace this portion of Zinc's electron notation with Argon's chemical symbol in brackets ([Ar].)
    • So, Zinc's electron configuration written in shorthand is [Ar]4s2 3d10.

Using an ADOMAH Periodic Table

  1. Understand the ADOMAH Periodic Table. This method of writing electron configurations doesn't require memorization. However, it does require a rearranged periodic table, because in traditional periodic table, beginning with fourth row, period numbers do not correspond to the electron shells. Find an ADOMAH Periodic Table, a special type of periodic table designed by scientist Valery Tsimmerman. It's easily found via a quick online search.[2]
    • In the ADOMAH Periodic Table, horizontal rows represent groups of elements, such as halogens, inert gases, alkali metals, alkaline earths, etc. Vertical columns correspond to electron shells and so called “cascades” (diagonal lines connecting s,p,d and f blocks) correspond to periods.
    • Helium is moved next to Hydrogen, since both of them are characterized by the 1s orbital. Blocks of periods (s,p,d and f) are shown on the right side and shell numbers are shown at the base. Elements are presented in rectangular boxes that are numbered from 1 to 120. These numbers are normal atomic numbers that represent total number of electrons in a neutral atom.
  2. Find your atom in the ADOMAH table. To write electron configuration of an element, locate its symbol in ADOMAH Periodic Table and cross out all elements that have higher atomic numbers. For example, if you need to write electron configuration of Erbium (68), cross out elements 69 through 120.
    • Notice numbers 1 through 8 at the base of the table. These are electron shell numbers, or column numbers. Ignore columns which contain only crossed out elements. For Erbium remaining columns are 1,2,3,4,5 and 6.
  3. Count orbital sets up to your atom. Looking at the block symbols shown on the right side of the table (s, p, d, and f) and at the column numbers shown at the base and ignoring diagonal lines between the blocks, break up columns into column-blocks and list them in order from the bottom up. Again, ignore column blocks where all elements are crossed out. Write down the column-blocks beginning with the column number followed by the block symbol, like this: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (in case of Erbium).
    • Note: The above electron configuration of Er is written in the order of ascending shell numbers. It could also be written in the order of orbital filling. Just follow cascades from top to bottom instead of columns when you write down the column-blocks: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f12.
  4. Count electrons for each orbital set. Count elements that were not crossed out in each block-column, assigning one electron per element, and write down their quantity next to the block symbols for each block-column, like this: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f12 5s2 5p6 6s2. In our example, this is is the electron configuration of Erbium.
  5. Know irregular electron configurations. There are eighteen common exceptions to electron configurations for atoms in the lowest energy state, also called the ground state. They deviate from the general rule only by last two-to-three electron positions. In these cases, the actual electron configuration keeps the electrons in a lower-energy state than in a standard configuration for the atom. The irregular atoms are:
    • Cr (..., 3d5, 4s1); Cu (..., 3d10, 4s1); Nb (..., 4d4, 5s1); Mo (..., 4d5, 5s1); Ru (..., 4d7, 5s1); Rh (..., 4d8, 5s1); Pd (..., 4d10, 5s0); Ag (..., 4d10, 5s1); La (..., 5d1, 6s2); Ce (..., 4f1, 5d1, 6s2); Gd (..., 4f7, 5d1, 6s2); Au (..., 5d10, 6s1); Ac (..., 6d1, 7s2); Th (..., 6d2, 7s2); Pa (..., 5f2, 6d1, 7s2); U (..., 5f3, 6d1, 7s2); Np (..., 5f4, 6d1, 7s2) and Cm (..., 5f7, 6d1, 7s2).

Tips

  • When the atom is an ion, it means that the number of protons does not equal the number of electrons. The charge of the atom will them be displayed at the top right (usually) corner of the chemical symbol. So, an antimony atom with charge +2 has an electron configuration of 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1. Notice that the 5p3 changed into a 5p1. Be careful when the configuration of an uncharged atom ends in anything but an s and p orbital set. When you take away electrons, you can only take them away from the valence orbitals (the s and p orbitals). So, if a configuration ends in 4s2 3d7, and the atom gains a charge of +2, then the configuration would change to end with 4s0 3d7. Notice that 3d7 does not change, instead, the s orbital electrons are lost.
  • Every atom desires to be stable, and the most stable configurations have full s and p (s2 and p6) orbital sets. The noble gases have this configuration, which is why they are rarely reactive and are on the right side of the periodic table. So, if a configuration ends in 3p4, it only needs two more electrons to become stable (losing six, including the s orbital set's electrons, takes more energy, so losing four is easier). And if a configuration ends in 4d3, it only needs to lose three electrons to reach a stable state. Also, half filled shells (s1, p3, d5..) are more stable than, for example p4 or p2; however, s2 and p6 will be even more stable.
  • You can also write an element's electron configuration by just writing the valence configuration, which is the last s and p orbital set. So, the valence configuration of an antimony atom would be 5s2 5p3.
  • Ions aren't the same. They're much harder. Skip two above levels of this article and follow the same pattern depending on where you started depending on how high or how low the number of electrons is.
  • To find the atomic number of the atom when it is in electron configuration form, just add up all of the numbers that follow the letters (s, p, d, and f). This only works if this is a neutral atom, if it is an ion, this does not work, and you would have to add or subtract however many electrons were added or lost.
  • The number following the letter is actually superscript, so do not make that mistake on a test.
  • There are two different ways to write electron configurations. They can be written in the order of ascending shell numbers, or in the order of orbital filling, as presented above for Erbium.
  • There are circumstances when an electron needs to be "promoted." When an orbital set is one electron away from being half occupied or completely occupied, remove one electron from the nearest s or p orbital set and move it to the orbital set that needs the electron.
  • There is no such thing as the "stability of a half-filled" sublevel. It is an oversimplification. Any stability relating to "half-filled" sub-levels is due to the fact that each orbital is singly occupied, thus electron-electron repulsions are minimized.

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Sources and Citations

  1. ↑ https://www.mnemonic-device.com/chemistry/sober-physicists-dont-find-giraffes-hiding-in-kitchens/
  2. ↑ http://www.meta-synthesis.com/webbook/35_pt/pt_database.php?PT_id=32
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