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Chemistry LibreTexts

Connecting Electronic Configurations to the Periodic Table

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  • Electron configuration can be described as how electrons are assembled within the orbitals shells and subshells of an atom. It is important to understand what an electron is in order to fully understand the electron configuration. An electron is a sub atomic particle that is associated with a negative charge. Electrons are found outside of the nucleus, as opposed to neutrons (particles with neutral charge,) and protons (particles with positive charge.) Furthermore, electrons are associated with energy, more specifically quantum energy, and exemplify wave-like and particle-like characteristics. The word configuration simply means the arrangement of something. Therefore electron configuration in straightforward language means the arrangement of electrons.


    In general when filling up the electron diagram, it is customary to fill the lowest energies first and work your way up to the higher energies. Principles and rules such as the Pauli exclusion principle, Hund’s rule, and the Aufbau process are used to determine how to properly configure electrons. The Pauli exclusion rule basically says that at most, 2 electrons are allowed to be in the same orbital. Hund’s rule explains that each orbital in the subshell must be occupied with one single electron first before two electrons can be in the same orbital. Lastly, the Aufbau process describes the process of adding electron configuration to each individualized element in the periodic table. Fully understanding the principles relating to electron configuration will promote a better understanding of how to design them and give us a better understanding of each element in the periodic table. How the periodic table was formed has an intimate correlation with electron configuration. After studying the relationship between electron configuration and the period table, it was pointed out by Niels Bohr that electron configurations are similar for elements within the same group in the periodic table. Groups occupy the vertical rows as opposed to a period which is the horizontal rows in the table of elements.

    S, P, D, and F Blocks

    • S block: The S block in the periodic table of elements occupies the alkali metals and alkaline earth metals, also known as groups 1 and 2. Helium is also part of the S block. The principal quantum number “n” fills the s orbital. There is a maximum of two electrons that can occupy the s orbital.
    • P Block: The P block contains groups 13, 14, 15, 16, 17, and 18, with the exception of Helium. (Helium is part of the S block.) The principal quantum number “n” fills the p orbital. There is a maximum of six electrons that can occupy the p orbital.
    • D Block: The D block elements are found in groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12 of the periodic table. The D block elements are also known as the transition metals. The d orbital is filled with the electronic shell “n-1.” There is a maximum of ten electrons that can occupy the d orbital.
    • F Block: The F block elements are the lanthanides and actinides. The F orbitals of the electron shell are filled with “n-2.” There is a maximum of fourteen electrons that can occupy the f orbital.

    Writing Electron Configuration in Noble Gas Notation

    Writing electron configurations of an element in noble gas notation makes electron configuration a much easier task. Noble Gas notation helps simplify writing out electron configurations and the concise form makes locating an element when given the electron configuration easier.

    Example 1: Oxygen

    Write the electron configuration of the element oxygen atoms.


    Oxygen has atomic number 8 in the periodic table and is located in period 2 and group 17 of the periodic table of elements. In order to write the electron configuration of this element, we use the noble gas preceding the period. In the case of Oxygen, the noble gas we use is Helium (He). We begin the configuration with the noble gas in brackets: [He]. We then continue writing the electron configuration from the noble gas (He) to the element oxygen as we normally would. Oxygen then becomes [He]2s22p4. When writing the configuration for elements farther down the periodic table, such as barium or iodine, this notation is especially helpful.

    Similarities Within A Group

    It is easy to see how similar electron configurations are in a group when written out. (Allow “n” to be the principal quantum number.) Lets first take a look at group 1 atoms. Group 1 atoms are the alkali metals. Let n=1. Notice the similar configuration within all the group 1 elements.

    Group Element Configuration
    1 H 1s1
    1 Li [He]2s1
    1 Na [Ne]3s1
    1 K [Ar]4s1
    1 Rb [Kr]5s1
    1 Cs [Xe]6s1
    1 Fr [Rn]7s1

    Now consider group 16 elements. These elements also will also have similar electron configurations to each another because they are in the same group; these elements have 6 valence electrons.

    Group Element Configuration
    16 O [He]2s22p4
    16 S [Ne]3s23p4
    16 Se [Ar]3d104s2 4p4
    16 Te [Kr]4d105s2 5p4
    16 Po [Xe]4f14 5d106s2 6p4

    Outside links


    1. Petrucci, Ralph H et al. General Chemistry: Principles & Modern Applications Ninth Edition. , Upper Saddle River, NJ: Pearson Prentice Hall, 2007. Print.
    2. Science Daily. 11/29/2009. <> Science Online Journal.
    3. Miloslav, Nic. International Union of Pure and Applied Chemistry. 2005-2009. "IUPAC Gold Book. " Online Book.

    Outside Links


    1. Question, True or False: Elements in the same period have similar electron configurations.

    Answer: False. Elements in the same GROUP have similar electron configurations.

    2. Question: What element has the electron configuration [Ar] 4s2 3d10 4p5?

    Answer: Bromine

    3. Question: What element has the electron configuration [Xe] 4f14 5d10 6s2 6p3?

    Answer: Bismuth

    4. Question: Demonstrate how elements in a group share similar characteristics by filling in the electron configurations for the Group 18 elements:

    Group Element Configuration
    18 He  
    18 Ne  
    18 Ar  
    18 Kr  
    18 Xe  
    18 Rn  

    Answer: 1s2, [He]2s22p6, [Ne]3s23p6, [Ar]3d104s24p6, [Kr]4d105s25p6, [Xe]4f145d106s26p6

    5. Question: How many valence electrons are there in Iodine?

    Answer: Iodine, z=53, group 17. This means there are seven valence electrons.

    6. Question: What is the highest number of electrons a 4p subshell can hold?

    Answer: 6! Each 3 p orbital can hold 2 electrons so if they are all filled, the answer is 6. You get this by multiplying the three orbitals by 2 electrons per orbital, so 3 multiplied by 2 equals 6.

    Make up some practice problems for the future readers.


    • Mariana Gerontides (UCD)