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• To step
• Method 1 of 2: Method one: Allocate electrons using the periodic table
• Method 2 of 2: Method Two: Use an ADOMAH periodic table
• Tips

The electron configuration of an atom is a numerical representation of the electron orbitals. Electron orbitals are differently shaped areas around the nucleus of an atom, where it is mathematically demonstrable that there is a chance that electrons exist there. An electron configuration makes it easy and quick to read how many electron orbitals an atom has and how many electrons are present in each orbital. Here you will learn how to start making your own electron configuration.

To step

Method 1 of 2: Method one: Allocate electrons using the periodic table

1. Find the atomic number. Each atom has an associated specific number of electrons. Find your atom's chemical symbol in the periodic table. The atomic number is a positive integer that starts at 1 (for hydrogen) and increases by 1 for each subsequent atom. The atomic number is the number of protons in that atom - so it is also the number of electrons in that atom if it is uncharged.
2. Determine the charge of the atom. Uncharged atoms have exactly the same number of protons as electrons, as indicated in the periodic table. But this is not the case with charged atoms. If you are dealing with a charged atom, add or subtract the electrons as follows: add one electron for each negative charge and subtract one for each positive charge.
   • For example: A sodium atom with a charge of -1 then has one extra electron added to the atomic number of 11. So this sodium atom has 12 electrons in total.
3. Memorize the basic list of orbitals. When an atom gains electrons, they fill different series of orbitals in a fixed order. Each orbital, when full, contains a fixed number of electrons. The orbital shapes are:
   • The s orbital (each number in the electron configuration followed by an "s") contains a single orbital, and through it Exclusion principle of Pauli we know that a single orbital can hold a maximum of 2 electrons, so each orbital shape can hold 2 electrons.
   • The p orbital contains 3 orbitals, so it can hold a total of 6 electrons.
   • The d orbital contains 5 orbitals, so it can hold 10 electrons.
   • The f orbital contains 7 orbitals, so it can hold 14 electrons.
4. Understand the notation of an electron configuration. Electron configurations are noted in such a way that it is clear how many electrons are present in the atom, and how many electrons are in each orbital. An orbital has a fixed notation with the number of electrons in superscript after the name of the oribital. The final electron configuration is a series of orbital shapes and superscripts.
   • For example, a simple electron configuration: 1s 2s 2p. This configuration indicates that there are two electrons in the 1s orbital shape, two electrons in the 2s orbital shape, and six electrons in the 2p orbital shape. 2 + 2 + 6 = 10 electrons in total. This is the electron configuration of an uncharged neon atom (Ne; atomic number 10.)
5. Learn the order of the orbitals. Note that the orbital shapes are numbered according to the electron shell, but ordered by energy level. For example, a fully filled 4s has less energy (or less potential) than a partially filled or filled 3d, so the 4s shell is in front. If you know the order of the orbitals, it is not difficult to fill them according to the number of electrons in the atom. The order in which the orbitals are filled is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
   • An electron configuration of an atom in which each orbital is completely filled is noted as follows: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d7p
   • Note that in the above list, if all the shells are filled, it is the electron configuration of ununoctium (Uuo; atomic number 118), the highest number in the periodic table - so this electron configuration contains every electron shell now known in an uncharged atom.
6. Fill the orbitals according to the number of electrons in your atom. For example, if we want to write down the electron configuration of an uncharged calcium atom, we start by looking up the atomic number in the periodic table. The atomic number of calcium is 20, so we write a configuration for this atom with 20 electrons in the order as indicated above.
   • Fill the orbitals with electrons according to the order above until you reach twenty. 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. So, the electron configuration for calcium is: 1s 2s 2p 3s 3p 4s.
   • Note: Energy levels change as you level up. For example, if you are about to move on to the 4th energy level, it first becomes 4s, afterwards 3d. After the fourth level, you move on to the fifth level, where the normal order is maintained. This only happens after the 3rd energy level.
7. Use the periodic table as a visual aid. You may have noticed that the order of the periodic table corresponds to the order of the orbital shapes in electron configurations. For example, atoms in the second column on the left always end with "s", atoms on the far right in the narrow middle portion always end with "d," etc. Use the periodic table as a visual guide to note configurations - the order in which your electrons adding to the orbitals corresponds to the position in the table of the periodic table. Take a good look at the following:
   • The two columns on the far left are a representation of atoms whose electron configurations end in s orbitals, the right block of this table is a representation of atoms whose configurations end in p orbitals, the central part, the atoms ending in a d orbital , and the lower region, atoms ending in f orbitals.
   • For example, when you write down an electron configuration for chlorine (Cl), remember, "This atom is in the third row (or" period ") of the periodic table. It is also in the fifth column of the p-orbitals group. this electron configuration ends in ... 3p
   • Note - the groups d and f orbitals in the table correspond to energy levels that are different from the period in which they are located. For example, the first row of the group of d orbitals corresponds to the 3d orbital even though it is in period 4, while the first row of the f orbitals corresponds to the 4f orbital even though it is in the sixth period.
8. Learn a shorthand for writing long electron configurations. The atoms along the right side of the periodic table are called the noble gases . These elements are very stable. To shorten the process of notating a long electron configuration, write the chemical symbol of the nearest gas, with fewer electrons than your atom, in square brackets, then continue with the electron configuration for the following orbital shapes. See below:
   • To understand this concept it is useful to write down an example of a configuration. Let's write the configuration of zinc (atomic number 30) using the abbreviated notation for a noble gas. The full electron configuration of zinc is: 1s 2s 2p 3s 3p 4s 3d. But note that 1s 2s 2p 3s 3p is the configuration of the noble gas argon. Simply replace this part of the zinc notation with the chemical symbol of argon in square brackets ([Ar].)
   • So the shorthand notation of the electron configuration of zinc can be written as [Ar] 4s 3d.

Method 2 of 2: Method Two: Use an ADOMAH periodic table

1. Understanding the ADOMAH Periodic Table. This method of noting electron configurations does not require much memorization. But it does require a different periodic table, because within the traditional periodic table, the electron shells, starting from the fourth row, do not correspond to the periodic numbers. Try to find an example of this system designed by Valery Tsimmerman online. Surely this is not a problem.
   • Within the ADOMAH Periodic Table, the rows represent groups of elements such as halogens, inert gases, alkali metals, etc. The columns correspond to the electron shells and the “cascades” (diagonal lines connecting s, p, d and f groups) correspond to the periods.
   • Helium now stands next to hydrogen because both are characterized by the 1s orbital. The periods (s, p, d and f) are on the right and the shell numbers at the bottom of the table. Elements are displayed in boxes numbered 1 through 120. These numbers represent the ordinary atomic numbers and indicate the number of electrons in a neutral atom.
2. Look for your atom in the ADOMAH table. To write down the electron configuration of an element, search for the appropriate symbol in the ADOMAH Periodic Table and cross through all elements with higher atomic numbers. For example, if you want to know the electron configuration of erbium (68), cross through elements 69 to 120.
   • Go to numbers 1 through 8 at the bottom (the base) of the table. These are the numbers of the electron shells, or the columns. Ignore the columns with crossed out elements. The remaining columns for erbium are 1,2,3,4,5 and 6.
3. Count the orbitals up to your atom. By looking at the group of symbols on the right side of the table (s, p, d, and f) and the column numbers at the bottom of the table and by ignoring the diagonal lines between them, you can divide the columns into groups and these in a list from bottom to top. Again, ignore those blocks with all elements crossed out. Write down the column groups, starting with the column number followed by the group symbol, like this: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (in the case of erbium).
   • Note: The above electron configuration of Er (erbium) is listed in the order of ascending shell numbers. It can also be written in the order of the orbitals. Just follow the cascades from top to bottom, instead of the columns, when you are writing the column groups: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f.
4. Count the electrons of each orbital shape. Count the elements that are not crossed out in each column group, picking one electron per element, and write the number next to the group symbols of each column group, like this: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s. This is the electron configuration of Er (erbium) from our example.
5. Know the irregular electron configurations. There are eighteen exceptions to the electron configurations of atoms of the lowest energy level, also known as the ground state. These deviate from the general rule for the last two or three electron positions. In these cases, the actual electron configurations keep the electrons at a lower energy level than a standard configuration of that 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); Dad (..., 5f2, 6d1, 7s2); YOU (..., 5f3, 6d1, 7s2); Np (..., 5f4, 6d1, 7s2) and Cm (..., 5f7, 6d1, 7s2).

Tips

• To find the atomic number of an atom if it is written in the form of an electron configuration, add all the numbers that come after the letters (s, p, d, and f). This only works with a neutral atom, not an ion, and you have to subtract or add any electrons that have disappeared or added.
• The number after the letter is actually in superscript, so make no mistake about this with a test.
• There is no such thing as the "stability of a half-filled" sub-level. This has been put too simply. The stability is because each orbital is occupied by only one electron, so electron-electron repulsion is minimal.
• Every atom wants to return to a stable state, and the most stable configurations have completely filled s and p (s2 and p6) orbitals. The noble gases have this configuration, which is why they are almost never reactive and are located on the right side of the periodic table. So, if a configuration ends with 3p, it only needs two more electrons to become stable (losing six electrons, including that of the s orbital, requires more energy, so it is easier to lose four). And if a configuration ends with 4d, it only needs to lose three more electrons to get into a stable state. It also holds that half-filled shells (s1, p3, d5 ..) are more stable than, for example, p4 or p2; s2 and p6 will become even more stable.
• When the atom is an ion, it means that the number of protons is not equal to the number of electrons. The charge of the atom is usually indicated in the top right corner of the symbol. So, an antimony atom with a +2 charge has an electron configuration of 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p. Note that the 5p has changed to 5p. Be aware of this if the configuration of an uncharged atom ends on anything but an s and p orbital. If you remove electrons, you can only do this with the valence orbitals (the s and p orbitals). So if a configuration ends with 4s 3d, and the charge of the atom increases with +2, the configuration changes so that it ends with 4s 3d. Remember that 3dnot changes, but the s orbital loses its electrons.
• There are circumstances where an electron gets a higher level. When an orbital is only one electron away from being half or fully filled, remove an electron from the nearest s or p orbital and move it to the orbital that needs that electron.
• You can also write down the electron configuration of an element by simply writing down the valence configuration, the last s and p orbital. So, the valence configuration of antimony then becomes 5s 5p.
• Ions are not the same, but much more difficult. Skip two levels and then follow the same pattern depending on where you started and the number of electrons.