Content

• Steps
• Method 1 of 2: Distribution of electrons using the periodic system of D. I. Mendeleev
• Method 2 of 2: Using the ADOMAH Periodic Table
• Tips

Electronic configuration an atom is a numerical representation of its electron orbitals. Electronic orbitals are regions of various shapes located around an atomic nucleus in which an electron is mathematically probable. Electronic configuration helps to quickly and easily tell the reader how many electron orbitals an atom has, as well as determine the number of electrons in each orbital. After reading this article, you will have mastered the method of generating electronic configurations.

Steps

Method 1 of 2: Distribution of electrons using the periodic system of D. I. Mendeleev

1. 1 Find the atomic number of your atom. Each atom has a specific number of electrons associated with it. Find the symbol for your atom in the periodic table. An atomic number is a positive integer starting at 1 (for hydrogen) and increasing by one for each subsequent atom. An atomic number is the number of protons in an atom, and therefore it is also the number of electrons in an atom with zero charge.
2. 2 Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown in the periodic table. However, charged atoms will have more or fewer electrons, depending on the amount of their charge. If you are working with a charged atom, add or subtract electrons as follows: add one electron for each negative charge and subtract one for each positive one.
   • For example, a sodium atom with a charge of -1 will have an extra electron in addition to its base atomic number 11. In other words, the total atom will have 12 electrons.
   • If we are talking about a sodium atom with a charge of +1, one electron must be subtracted from the base atomic number 11. Thus, the atom will have 10 electrons.
3. 3 Remember the basic list of orbitals. As the number of electrons increases, they fill the various sublevels of the electron shell of the atom according to a certain sequence. Each sublevel of the electron shell, when filled, contains an even number of electrons. The following sublevels are available:
   • s-sublevel (any number in the electronic configuration that comes before the letter "s") contains a single orbital, and, according to Pauli's principle, one orbital can contain a maximum of 2 electrons, therefore, there can be 2 electrons on each s-sublevel of the electron shell.
   • p-sublevel contains 3 orbitals, and therefore can contain a maximum of 6 electrons.
   • d-sublevel contains 5 orbitals, so it can have up to 10 electrons.
   • f-sublevel contains 7 orbitals, so it can have up to 14 electrons.
   • g-, h-, i- and k-sublevels are theoretical. The atoms containing electrons in these orbitals are unknown. The g-sublevel contains 9 orbitals, so theoretically it could have 18 electrons. The h-sublevel may have 11 orbitals and a maximum of 22 electrons; in the i-sublevel -13 orbitals and a maximum of 26 electrons; in the k-sublevel - 15 orbitals and a maximum of 30 electrons.
   • Memorize the order of the orbitals using the mnemonic trick:
     Sober Physicists Don’t Find Giraffes Hiding In Kitchens (sober physicists don't find giraffes hiding in kitchens).
4. 4 Understand the electronic configuration record. Electronic configurations are recorded to clearly reflect the number of electrons in each orbital. Orbitals are written sequentially, with the number of atoms in each orbital being superscript to the right of the orbital name. The completed electronic configuration takes the form of a sequence of sublevel designations and superscripts.
   • For example, the simplest electronic configuration: 1s 2s 2p. This configuration shows that there are two electrons at the 1s sublevel, two electrons at the 2s sublevel, and six electrons at the 2p sublevel. 2 + 2 + 6 = 10 electrons in total. This is the electronic configuration of a neutral neon atom (neon atomic number is 10).
5. 5 Remember the order of the orbitals. Keep in mind that the electron orbitals are numbered in ascending order of the electron shell number, but in ascending order of energy. For example, a filled 4s orbital is less energetic (or less mobile) than a partially filled or filled 3d, so the 4s orbital is recorded first. Once you know the order of the orbitals, you can easily fill them in according to the number of electrons in the atom. The order of filling the orbitals is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
   • The electronic configuration of an atom in which all orbitals are filled will have the following form: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d7p
   • Note that the above entry, when all orbitals are filled, is the electronic configuration of the element Uuo (ununoctium) 118, the highest numbered atom in the periodic table. Therefore, this electronic configuration contains all currently known electronic sublevels of a neutral charged atom.
6. 6 Fill in the orbitals according to the number of electrons in your atom. For example, if we want to write down the electronic configuration of a neutral calcium atom, we must start by looking for its atomic number in the periodic table. Its atomic number is 20, so we will write the configuration of an atom with 20 electrons according to the above order.
   • Fill in the orbitals in the order above until you reach the twentieth electron. The first 1s orbital will contain two electrons, the 2s orbitals will also have two, 2p - six, 3s - two, 3p - 6, and 4s - 2 (2 + 2 + 6 +2 + 6 + 2 = 20 .) In other words, the electronic configuration of calcium is: 1s 2s 2p 3s 3p 4s.
   • Note that the orbitals are in ascending order of energy. For example, when you are ready to move to the 4th energy level, then first write down the 4s orbital, and then 3d. After the fourth energy level, you go to the fifth, where the same order is repeated. This happens only after the third energy level.
7. 7 Use the periodic table as a visual clue. You've probably already noticed that the shape of the periodic table corresponds to the order of electronic sublevels in electronic configurations. For example, the atoms in the second column from the left always end in "s", while the atoms on the right edge of the thin middle section always end in "d", and so on. Use the periodic table as a visual guide to writing configurations - as the order in which you add to orbitals corresponds to your position in the table. See below:
   • In particular, the two leftmost columns contain atoms whose electronic configurations end in s-orbitals, the right block of the table contains atoms whose configurations end in p-orbitals, and in the lower part, atoms end in f-orbitals.
   • For example, when you write down the electronic configuration of chlorine, think like this: "This atom is located in the third row (or" period ") of the periodic table. It is also located in the fifth group of the p orbital block of the periodic system. Therefore, its electronic configuration will end in. ..3p
   • Please note: the elements in the region of the d and f orbitals of the table are characterized by energy levels that do not correspond to the period in which they are located. For example, the first row of the block of elements with d-orbitals corresponds to 3d orbitals, although it is located in the 4th period, and the first row of elements with f-orbitals corresponds to the 4f orbital, despite the fact that it is in the 6th period.
8. 8 Learn the shorthand for writing long electronic configurations. The atoms on the right edge of the periodic table are called noble gases. These elements are chemically very stable. To shorten the process of writing long electronic configurations, simply write in square brackets the chemical symbol of the nearest noble gas with fewer electrons than your atom, and then continue writing the electronic configuration of subsequent orbital levels. See below:
   • To understand this concept, it is helpful to write an example configuration. Let's write the configuration for zinc (atomic number 30) using the noble gas abbreviation. The complete zinc configuration looks like this: 1s 2s 2p 3s 3p 4s 3d. However, we see that 1s 2s 2p 3s 3p is the electronic configuration of argon, a noble gas. Simply replace the electronic configuration portion of zinc with the chemical symbol argon in square brackets ([Ar].)
   • So, the electronic configuration of zinc, written in an abbreviated form, is: [Ar] 4s 3d.
   • Note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write [Ar]! One must use the reduction of the noble gas facing this element; for argon it will be neon ([Ne]).

Method 2 of 2: Using the ADOMAH Periodic Table

1. 1 Learn the ADOMAH periodic table. This method of recording the electronic configuration does not require memorization, however, it requires a revised periodic table, since in the traditional periodic table, starting from the fourth period, the period number does not correspond to the electron shell. Find the ADOMAH Periodic Table - a special type of periodic table developed by scientist Valery Zimmerman. It is easy to find it with a short internet search.
   • In the periodic table of ADOMAH, horizontal rows represent groups of elements such as halogens, noble gases, alkali metals, alkaline earth metals, etc. Vertical columns correspond to electronic levels, and so-called "cascades" (diagonal lines connecting blocks s, p, d and f) correspond to periods.
   • Helium is moved to hydrogen as both of these elements have a 1s orbital. Period blocks (s, p, d and f) are shown on the right side, and level numbers are shown at the bottom. Elements are shown in boxes numbered 1 through 120. These numbers are common atomic numbers that represent the total number of electrons in a neutral atom.
2. 2 Find your atom in the ADOMAH table. To record the electronic configuration of an element, find its symbol in the ADOMAH periodic table and cross out all elements with a higher atomic number. For example, if you need to write down the electronic configuration of erbium (68), cross out all elements from 69 to 120.
   • Note the numbers 1 through 8 at the bottom of the table. These are electronic level numbers, or column numbers. Ignore columns that only contain crossed out items.For erbium, the columns numbered 1,2,3,4,5 and 6 remain.
3. 3 Count the orbital sublevels to your element. Looking at the block symbols shown to the right of the table (s, p, d, and f) and the column numbers shown at the bottom, ignore the diagonal lines between the blocks and break the columns into column blocks in order from bottom to top. Again, ignore the boxes with all the elements crossed out. Write down the column blocks, starting with the column number followed by the block symbol, thus: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (for erbium).
   • Note: The above electronic configuration Er is written in ascending order of the electronic sublevel number. It can also be written in the order of filling the orbitals. To do this, follow the cascades from the bottom up, not the columns, when you write the column blocks: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f.
4. 4 Count the electrons for each electronic sublevel. Count the elements in each block-column that were not crossed out, attaching one electron from each element, and write their number next to the block symbol for each block-column as follows: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s ... In our example, this is the electronic configuration of erbium.
5. 5 Consider incorrect electronic configurations. There are eighteen typical exceptions related to the electronic configurations of atoms in the lowest energy state, also called the ground energy state. They do not obey the general rule only in the last two or three positions occupied by electrons. In this case, the actual electronic configuration assumes that the electrons are in a state with a lower energy in comparison with the standard configuration of the atom. Exception atoms include:
   • 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

• To find the atomic number of an atom when written in electronic configuration, simply add up all the numbers that follow the letters (s, p, d, and f). This only works for neutral atoms, if you are dealing with an ion, then nothing will work - you have to add or subtract the number of extra or lost electrons.
• The number following the letter is a superscript, do not make a mistake in the check.
• There is no "stability of a half-filled" sublevel. This is a simplification. Any stability that relates to the "half filled" sublevels is due to the fact that each orbital is occupied by one electron, so the repulsion between the electrons is minimized.
• Each atom tends to a stable state, and the most stable configurations have filled sublevels s and p (s2 and p6). Noble gases have such a configuration, therefore they rarely enter into reactions and are located on the right in the periodic table. Therefore, if the configuration ends at 3p, then it needs two electrons to reach a stable state (to lose six, including electrons of the s-sublevel, it will take more energy, so it is easier to lose four). And if the configuration ends in 4d, then it needs to lose three electrons to reach a stable state. In addition, half-filled sublevels (s1, p3, d5 ..) are more stable than, for example, p4 or p2; however, s2 and p6 will be even more stable.
• When you are dealing with an ion, this means that the number of protons is not equal to the number of electrons. In this case, the charge of an atom will be shown at the top to the right (as a rule) of the chemical symbol. Therefore, an antimony atom with a charge of +2 has the electronic configuration 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p. Note that 5p has changed to 5p. Be careful when the configuration of a neutral atom ends up at sublevels other than s and p. When you pick up electrons, you can only pick them up from the valence orbitals (s and p orbitals).Therefore, if the configuration ends at 4s 3d and the atom gains a +2 charge, then the configuration will end at 4s 3d. Please note that 3d not changes, instead of losing s-orbital electrons.
• There are conditions when the electron is forced to "go to a higher energy level." When a sublevel lacks one electron to half or full filling, take one electron from the nearest s or p-sublevel and move it to the sublevel that needs an electron.
• There are two options for recording an electronic configuration. They can be written in ascending order of energy level numbers or in the order of filling of electron orbitals, as was shown above for erbium.
• You can also write down the electronic configuration of an element by writing down only the valence configuration, which is the last s and p sublevels. Thus, the valence configuration of antimony will have the form 5s 5p.
• Jonah is not the same. It is much more difficult with them. Skip two levels and follow the same pattern depending on where you started and how large the number of electrons is.