The electron configuration of an atom is a numerical representation of its orbitals. Orbitals have different shapes and positions with respect to the nucleus, and represent the area in which you have the greatest chance of detecting an electron. The electron configuration quickly indicates how many orbitals an atom has and the amount of electrons that "populate" each orbital. When you understand the basic principles underlying electronic configuration and are able to write it down, then you can take any chemistry exam with confidence.
Steps
Method 1 of 2: With the Periodic Table
Step 1. Find the atomic number
Each atom is associated with an atomic number which indicates the number of protons. The latter, in a neutral atom, is equal to the number of electrons. The atomic number is a positive integer, hydrogen has an atomic number equal to 1, and this value increases by one as you move to the right on the periodic table.
Step 2. Determine the charge of the atom
Neutral ones have a number of electrons equal to the atomic number, while charged atoms can have a greater or lesser quantity, depending on the power of the charge; then add or subtract the number of electrons depending on the charge: add one electron for each negative charge and subtract one electron for each positive charge.
For example, a sodium atom with a negative -1 charge will have an "extra" electron of atomic number 11, hence 12 electrons
Step 3. Memorize the basic list of orbitals
Once you know the order of the orbitals, it will be easy to complete them according to the number of electrons in an atom. The orbitals are:
- The group of s-type orbitals (any number followed by an "s") contains a single orbital; according to the Pauli exclusion principle, a single orbital can contain a maximum of 2 electrons. It follows that each s orbital can contain 2 electrons.
- The p-type group of orbitals contains 3 orbitals, so it may contain a total of 6 electrons.
- The group of orbitals of type d contains 5 orbitals, so it can contain 10 electrons.
- The group of f-type orbitals contains 7 orbitals, so it can contain 14 electrons.
Step 4. Understand the electronic configuration notation
It is written in such a way that both the number of electrons in the atom and the number of electrons in each orbital appear clearly. Each orbital is written according to a certain sequence and with the number of electrons following the name of the orbital itself. The final configuration is a single row of orbital and superscript names.
For example, here is a simple electronic configuration: 1s2 2s2 2p6. You can see that there are two electrons on the 1s orbital, two in the 2s orbital and 6 in the 2p orbital. 2 + 2 + 6 = 10 electrons in all. This configuration refers to a neutral neon atom (which has an atomic number of 10).
Step 5. Memorize the order of the orbitals
Remember that the groups of orbitals are numbered according to the electron shell, but ordered in terms of energy. For example, a full 4s orbital2 has a lower (or potentially less unstable) energy level than a partially full or completely full 3d one10; it follows that 4s will come first in the list. When you know the order of the orbitals you simply have to fill in the diagram with the number of electrons of the atom. The order 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 all orbitals occupied should be written like this: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d107p68s2.
- Note that the above example, if all the electronic shells were complete, would indicate the electronic configuration of the ununoctio (Uuo), 118, the atom with the largest atomic number in the periodic table of elements. This electronic configuration contains all known electronic shells for a neutral atom.
Step 6. Fill the orbitals according to the number of electrons in your atom
For example, let's write the electron configuration of a neutral calcium atom. First we need to identify the atomic number in the periodic table. This number is 20, so we need to write the electronic configuration of an atom with 20 electrons following the order described above.
- Fill the orbitals in order until you have placed all 20 electrons. The 1s orbital has two electrons, the 2s has two, the 2p has six, the 3s has six and the 4s has two (2 + 2 + 6 +2 +6 + 2 = 20). So the electron configuration for a neutral calcium atom is: 1s2 2s2 2p6 3s2 3p6 4s2.
- Note: The energy level varies as you move up into the orbitals. For example, when you are about to rise to the fourth energy level, first comes 4s, after 3d. After the fourth level, you will move on to the fifth level, which again follows the normal order. This only happens after the third energy level.
Step 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 the orbitals in an electron configuration. For example, the atoms in the second column from the left always end in "s2", the ones to the right of the narrowest central part always end in" d10", and so on. Then use the periodic table as a guide for writing the configuration; the order in which you add electrons to the orbitals corresponds to the position in the table. Here's how:
- Specifically, the two leftmost columns represent the atoms whose configuration ends with an s orbital, the block on the right of the table represents the atoms whose configuration ends with a p orbital, while the central section encloses the atoms that have a configuration ending with an orbital d. The lower portion of the periodic table contains atoms with a configuration ending in an f orbital.
- For example, if you have to write the electron configuration of chlorine, think: "this atom is in the third row (or" period ") of the periodic table. It is also in the fifth column so the configuration ends with … 3p5".
- Warning: the d and f orbitals of the elements of the periodic table have different energy levels compared to the period in which they are inserted. For example, the first row of the d-orbital block corresponds to the 3d orbital even though it is within period 4, while the first row of the f-orbital corresponds to 4f even though it is within period 6.
Step 8. Learn some tricks for writing long electronic configurations
The atoms at the right end of the periodic table are called noble gases. These are very stable elements. To shorten the writing of a long configuration, simply write, in square brackets, the chemical symbol of the noble gas with fewer electrons than the element you are considering, and then continue writing the configuration for the remaining electrons.
- An example is useful to understand the concept. We write the electron configuration of zinc (atomic number 30) using a noble gas as a shortcut. The full configuration for zinc is: 1s2 2s2 2p6 3s2 3p6 4s2 3d10. However, you may notice that 1s2 2s2 2p6 3s2 3p6 is the configuration of argon, a noble gas. So you can replace this part of the electron configuration of zinc with the argon symbol enclosed in square brackets ([Ar]).
- So you can write that the electron configuration of zinc is: [Ar] 4s2 3d10.
Method 2 of 2: With the ADOMAH Periodic Table
Step 1. To write the electronic configurations there is an alternative method that requires neither memorization nor mnemonic diagrams
However, it requires a modified periodic table. In the traditional one, from the fourth line, the periodic numbers do not correspond to the electronic shells. This special board was developed by Valery Tsimmerman and you can find it on the website: (www.perfectperiodictable.com/Images/Binder1).
- In the periodic table ADOMAH the horizontal lines represent the groups of elements, such as halogens, inert gases, alkali metals, alkaline earths, etc. The vertical columns correspond to the electronic shells and the so-called "cascades" correspond to the periods (where diagonal lines join the blocks s, p, d and f).
- Helium is found close to hydrogen, as they are both characterized by electrons located in the same orbital. The blocks of the periods (s, p, d and f) appear on the right, while the numbers of the shells are found at the bottom. The elements are represented in rectangles numbered from 1 to 120. These are called atomic numbers and also represent the total number of electrons in a neutral atom.
Step 2. Print a copy of the ADOMAH periodic table
To write the electronic configuration of an element, look for its symbol in the ADOMAH table, and delete all elements that have a higher atomic number. For example, if you have to write the electronic configuration of the erbium (68), delete the elements starting from 69 up to 120.
Consider the numbers 1 to 8 at the base of the table. These are the numbers of the electronic shells, or the numbers of the columns. Disregard columns in which all elements are deleted. Those that remain for erbium are 1, 2, 3, 4, 5 and 6
Step 3. Look at the block symbols on the right of the table (s, p, d, f) and the column numbers below; ignore the diagonal lines between the various blocks, separate the columns into column-block pairs and order them from bottom to top
Again, don't consider blocks where the elements are all deleted. Write the column-block pairs starting with the number of columns followed by the block symbol, as indicated here: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (in the case of erbium).
Note: the electronic configuration of the ER reported above is written in ascending order with respect to the number of shells. One could also write in the order of filling of the orbitals. Simply, you have to follow the cascades from top to bottom instead of columns when writing column-block pairs: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f12.
Step 4. Count the elements that are not deleted in each block-column and write this number next to the block symbol, as below:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f12 5s2 5p6 6s2. This is the electronic configuration of erbium.
Step 5. There are eighteen common exceptions to the electronic configurations of atoms in the lowest energy level, also referred to as the base state
They deviate from the general rule only in the penultimate and third to last position of the electrons. Here they 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); There(…, 5d1, 6s2); There is(…, 4f1, 5d1, 6s2); Gd(…, 4f7, 5d1, 6s2); Au(…, 5d10, 6s1); B. C(…, 6d1, 7s2); Th(…, 6d2, 7s2); Pa(…, 5f2, 6d1, 7s2); U(…, 5f3, 6d1, 7s2); Np(…, 5f4, 6d1, 7s2) e Cm(…, 5f7, 6d1, 7s2).
Advice
- To find the atomic number of an element, given the electronic configuration, add together all the numbers following the letters (s, p, d, and f). This only works if the atom is neutral; if you are dealing with an ion you have to add or subtract as many electrons based on the charge.
- The numbers following the letters are quotation marks, so don't get confused when checking.
- There is no such thing as "the stability of a half-filled sublevel". It is an oversimplification. Any stability that refers to a "half-completed" level is due to the fact that each orbital is occupied by a single electron and that electron-electron repulsion is minimal.
- When you have to work with an ion, it means that the number of protons is not equal to that of electrons. The charge is usually expressed at the top right of the chemical symbol. So an antimony atom with a +2 charge has an electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1. Note that 5p3 changed to 5p1. Be very careful when the configuration of a neutral atom ends in something other than an s and p orbital. When you take out electrons, you can't do it from valence orbitals (like s and p). So if the configuration ends with 4s2 3d7, and the atom has a +2 charge, then the configuration changes in 4s0 3d7. Note that 3d7Not changes; while the electrons of the s orbital are lost.
- Each atom tends towards stability, and the most stable configurations have complete s and p orbitals (s2 and p6). Noble gases have this configuration and are on the right side of the periodic table. So if the configuration ends with 3p4, it only takes two more electrons to become stable (losing six takes too much energy). And if the configuration ends with 4d3, it is enough to lose three electrons to achieve stability. Again, half-complete shells (s1, p3, d5..) are more stable than, for example, p4 or p2; however, s2 and p6 will be even more stable.
- There are two different ways of writing the electronic configuration: in ascending order of electronic shells or according to the order of orbitals, as written above for erbium.
- There are circumstances in which an electron has to be "promoted". When only one electron is missing in an orbital to be complete, remove an electron from the nearest s or p orbital and move it to the orbital that needs to be completed.
- You can also write the electronic configuration of an element simply by writing the valence configuration, i.e. of the last s and p orbitals. Hence the valence configuration of an antimony atom is 5s2 5p3.
- The same is not true for ions. Here the question becomes a little more difficult. The number of electrons and the point at which you started skipping the levels will determine the compilation of the electronic configuration.
