Assigning Electrons Using a Periodic Table
- 1Find 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 0 charge.2Determine the charge of the atom. Uncharged atoms will have exactly the number of electrons as is represented on the periodic table. However, charged atoms (ions) 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 1 electron for each negative charge and subtract 1 for each positive charge.
3Memorize 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:
- 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.
- A sodium atom with a -1 charge would have 1 electron added to its basic atomic number of 11. The sodium atom would then have a total of 12 electrons.
4Understand 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.
- 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.
- The g, h, i and k orbital sets are theoretical. No known atoms have electrons in any of these orbitals. The g set has 9 orbitals, so it could theoretically contain 18 electrons. The h set would have 11 orbitals and a maximum of 22 electrons, the i set would have 13 orbitals and a maximum of 26 electrons, and the k set would have 15 orbitals and a maximum of 30 electrons.
- Remember the order of the letters with this mnemonic:Sober Physicists Don’t Find Giraffes Hiding In Kitchens.
5Memorize 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.
- For example, here is a simple electron configuration: 1s2 2s2 2p6. This configuration shows that there are 2 electrons in the 1s orbital set, 2 electrons in the 2s orbital set, and 6 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.)
6Fill 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.
- 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.
7Use 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.
- Fill up orbitals according to the order above until you reach 20 total electrons. The 1s orbital gets 2 electrons, the 2s gets 2, the 2p gets 6, the 3s gets 2, 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 4th energy level, you’ll move onto the 5th where it follows the order once again (5s, then 4d). This only happens after the 3rd energy level.
8Learn 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 fewer electrons than your atom in brackets, then continue with the electron configuration for the following orbital sets.
- Specifically, the 2 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 from 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.
- 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.
- Note that if you are doing noble gas notation for, say, argon, you cannot write [Ar]! You have to use the noble gas that comes before that element; for argon, that would be neon ([Ne]).
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