Electron configuration
Electron configuration describes the distribution of electrons in an atom and is crucial for understanding the chemical behavior of elements. Atoms consist of a nucleus, containing protons and neutrons, surrounded by a cloud of electrons, which do not follow fixed orbits but rather exist in probabilistic regions determined by their energy levels. The arrangement of these electrons, particularly the valence electrons in the outermost shell, influences how an atom interacts with others, thereby affecting its reactivity and bonding characteristics.
The concept of electron configuration evolved from early models like the Bohr model, which suggested discrete energy levels for electrons. Today, the quantum mechanical model provides a more accurate framework, describing electrons as existing in shells and subshells, each identified by principal quantum numbers and letter designations (s, p, d, f) that indicate their shape and capacity.
For example, the electron configuration of lawrencium (element 103) illustrates the specific arrangement of its 103 electrons across various subshells. Understanding electron configuration is fundamental, especially in predicting chemical reactions and bonding, as atoms typically seek to attain stable configurations, often guided by principles such as the octet rule. The complexity increases with transition metals, where valence electrons can occupy d and f subshells, leading to diverse chemical properties.
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Electron configuration
Electron configuration refers to the arrangement of electrons within an atom. In chemistry, electron configurations play a large role in explaining how different chemicals interact. They can be used to predict whether certain elements or compounds will react and in what way.
In order to appreciate the significance of electron configurations, it is important to first understand the anatomy of an atom. Atoms are the building blocks of all matter. They consist of a nucleus, which contains subatomic particles called protons and neutrons, surrounded by a diffuse cloud of electrons. While these electrons do not orbit the nucleus as neatly as, for example, planets orbiting a star, and in fact the location of a given electron within an atom cannot be pinpointed with absolute certainty, the rules of electron configuration can be used to determine an electron’s most likely position and how it will interact with the electrons of other atoms.
Brief History
In 1913, Danish physicist Niels Bohr (1885–1962) proposed that the properties of the elements could be explained by their individual electron configurations. He introduced what is now known as the Bohr model of the atom, in which electrons orbited the nucleus in a manner similar to how planets orbit the sun. Electrons in this model followed certain orbital paths that corresponded with fixed energy states. An electron would have to absorb a single unit, or quantum, of energy to move to a higher orbital or release a quantum of energy to drop back down to a lower orbital. While Bohr’s model has since been displaced by the quantum mechanical model of the atom, it was the first atomic model to incorporate quantum theory and is credited with introducing the concept of quantized—that is, discrete, rather than continuous—energy levels in electrons.
Overview
Despite what scientists used to believe, electrons do not travel around the nucleus in a circular fashion. Instead, they vibrate back and forth in certain patterns based on their different energy levels. Electron configuration is important because it is an atom’s valence electrons—that is, the electrons in its outermost electron shell—that determine how the atom reacts with other atoms; only the valence electrons take part in chemical bond formation.
In the quantum mechanical model of the atom, electrons can be thought of as occupying concentric electron shells surrounding the nucleus. While this model is not strictly accurate, it is a useful way to visualize an extremely complicated concept. Each electron shell is divided into a certain number of subshells, which in turn contain a number of orbitals that can hold up to two electrons each. These orbitals are not fixed paths that the electrons follow around the nucleus, but rather descriptions of the areas within the atom in which the electrons are most likely to be found.
Each electron shell is identified by a number (n = 1, n = 2, and so on), known as the principal quantum number, which refers to the principal energy level of the electrons contained within. The numbering proceeds outward from the nucleus, with the n = 1 shell being the innermost. The subshells are identified by a letter—s, p, d, or f—that specifies both the number of orbitals in each subshell and the shape of each orbital within that subshell. An s subshell contains a single spherical orbital, which can hold up to two electrons; a p subshell contains three roughly dumbbell-shaped orbitals, which can hold two electrons each, for a total of six; a d subshell contains five orbitals total, four four-lobed X-shaped orbits and one shaped like a dumbbell with a torus around the middle, for an overall capacity of ten electrons; and an f subshell contains seven complicated orbital shapes that are difficult to describe, allowing it to contain fourteen electrons altogether. Not every electron shell contains all four subshells. The p subshell first appears in the n = 2 shell, the d subshell in the n = 3 shell, and the f subshell in the n = 4 shell. A theoretical g subshell is posited to exist, but no element has yet been found to contain one.
The electron configuration of a particular atom is given in the format 1s22s22p63s23p63d104s24p64d104f145s25p65d105f146s26p67s27p1. Specifically, this is the electron configuration of lawrencium (Lw), element 103 of the periodic table. The superscript number following each subshell designation represents the number of electrons present in that subshell. The atomic number of each element is equal to the number of protons in the nucleus of that element, and in a neutral atom (as opposed to an ion), the number of electrons is equal to the number of protons; thus, a lawrencium atom has 103 electrons, and accordingly, the superscript numbers add up to 103.
When two atoms form a covalent bond, they share either two, four, or six valence electrons (for a single, double, or triple bond) between them, so that the electron(s) of one atom occupy available space in the valence shell of the other. In main-group elements—that is, elements in groups (or columns) 1, 2, and 13 through 18 of the periodic table—electron configuration tends to follow what is known as the octet rule, in which atoms of different elements typically combine or otherwise react in such a way that each atom ends up with eight electrons in its valence shell. This is because main-group elements have their valence electrons in either the s subshell (groups 1–2) or the p subshell (groups 13–18), which between them hold a total of eight electrons. Hydrogen and helium are exceptions to this rule, as the n = 1 shell only contains an s subshell, so both elements can only hold a maximum of two electrons total.
For transition elements—a category that traditionally refers to the elements in groups 3 through 12 and also often includes the lanthanides and actinides, usually displayed in a separate block beneath the periodic table—the rules of electron configuration are more complicated. These elements have their valence electrons in the d or f subshells, so they may contain more than eight electrons in their outermost electron shell. In addition, in a transition element, the electron shell directly below the valence shell may not be completely filled. This is because electrons occupy the subshells in order of increasing energy level, and the energy levels of the subshells grow increasingly close together the farther they get from the nucleus, so that beyond the 3p subshell, they begin to overlap to a certain degree. As a result, instead of proceeding in a strictly alphanumeric order, the electrons fill the subshells in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p (and so on, in theory). Thus, in the lawrencium example above, the first five electron shells follow the expected subshell pattern, with the n = 4 and n = 5 shells containing full s, p, d, and f subshells. However, the n = 6 shell does not contain a d or an f subshell, yet the 7s subshell is filled and the 7p subshell contains a single electron. In such cases, electrons in both the outermost electron shell and the shell immediately below it—here, the n = 7 and n = 6 shells—can function as valence electrons, participating in the formation of atomic bonds.
Bibliography
Brown, Theodore L., et al. Chemistry: The Central Science. 12th ed. Upper Saddle River: Prentice, 2012. Print.
Chivers, Tristram. A Guide to Chalcogen-Nitrogen Chemistry. Hackensack World Scientific, 2005. Print.
De La Cruz, Jerri. "Exploring Electron Configuration: Understanding the Behavior of Atoms." Learnt, 12 Apr. 2023, learnt.io/blog/electron-configuration/. Accessed 12 Nov. 2024.
Ferrer-Vinent, Ignacio J. “Teaching SciFinder Basics to Organic Chemistry Students.” Science & Technology Libraries 31.2 (2012): 164–79. Print.
Francl, Michelle. “Back to Basics.” Nature Chemistry 1.9 (2009): 681. Print.
Gonick, Larry, and Craig Criddle. The Cartoon Guide to Chemistry. New York: Harper, 2005. Print.
Neuss, Geoffrey. "A Problem with Explaining the Electron Configuration of Scandium." Foundations of Chemistry, vol. 23, pp. 239-42, 20 Jan. 2021, doi.org/10.1007/s10698-020-09392-2. Accessed 28 Dec. 2022.
Shindy, H. A. “Basics, Mechanisms and Properties in the Chemistry of Cyanine Dyes: A Review Paper.” Mini-Reviews in Organic Chemistry 9.4 (2012): 352–60. Print.
Solomons, Graham, Craig Fryhle, and Scott Snyder. Organic Chemistry. 11th ed. Hoboken: Wiley, 2014. Print.
Spencer, James N., George M. Bodner, and Lyman H. Rickard. Chemistry: Structure & Dynamics. 5th ed. Hoboken: Wiley, 2012. Print.