Electrons in Atoms

Posted: August 25th, 2021

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Electrons in Atoms

Electrons are the smallest sub-atomic components that are negatively charged and always orbiting within a shell covering an atom nucleus (Manini 112). On the other hand, atoms are the basic units of matter, which hold the electrons arranged in the form of shells. The atoms build up into visible objects of nature, besides facilitating the understanding of the objectives.  Shells inside an atom are distributed in levels, arranged around the core part of an atom(Manini 127). Each is considered a level of energy. According to Manini, each of these levels of energy should be full of electrons before an electron moves to the next higher shell of energy. The quantity of electrons that are held under each shell differs from one energy level to the other (144). For instance, one shell can hold two or three electrons. The only elements with one and two electrons are hydrogen and helium respectively. The subsequent energy levels can have about eight electrons, eighteen and thirty-two, and so on.

Similarly, in the theory of electron collision, Burkeand Joachain (101-115) explain that the movement and arrangement of an electron in each shell or energy level are unique; that is, it is not a perfect ordinary circular motion experienced around. Moreover, the shells are divided further into subshells, which are energy levels embedded in the primary shell energy levels (118). The subshells are usually denoted by letters such as p, s, f, d, which represents particular electron quantity. For instance, an s subshell has an amount of two electrons, while a p subshell contains about six electrons.

How Are Electrons Organized Inside An Atom?

As revealed earlier in the introduction, electrons are held within a shell inside an atom in the form of energy levels, which is described in an energetic state or quantum numbers. According to the Pauli Exclusion Principle, it is only one atom per four quantum numbers, which limits the electron quantity existing in either a subshell or a shell (Krane 143). Typically, electrons are arranged in an atom following the ascending order of energy levels. For example, H, which has single electrons, will often be positioned in n =1 shell that contains an s amount of subshells in a single orbital. Conventionally, this becomes easy when labeling the combination of shell-subshell using the shell number and letter representing the particular shell. Typically, specifying quantum numbers is not necessary; however, since the only possible value in H atom is m_l = 0, it becomes necessary. On the part of helium, He atom, there are two electrons. Here, it is only possible for the second electron to enter 1s shell-subshell if the atoms spinning quantum number varies from that of the first electron. In this case, the two electrons quantum set will be described as {1, 0, 0, +1/2} and {1, 0, 0, -1/2}. Notably, therefore, the description indicated in the overall set complies with the Pauli Exclusion Principle.

            Lithium, Li is the next atom, which contains three electrons. Following the Pauli Exclusion Principle, the electrons should not be placed in the 1s shell-subshell since the third electron will usually have four quantum similar to their other two electrons, hence going against the principle. Therefore, the third electrons must be assigned to the next energy level combination. Besides, it is critical noting that the first energy level, ℓ=0 thus, cannot hold the third electron. Accordingly, it demands that the third electrons should be assigned to n =2 shell characterized by s (ℓ=0) sub-shell with p (ℓ=1) subshell. More so, the arrangement of electrons should be in ascending order, from the lowest, which implies that this electron should be positioned in 2s shell-subshell. As a result, this offers the required combination in line with the exclusion principle. The principle’s net effect is to limit the amount of electrons assigned to the shell-subshell combination. If an s subshell is filled up, any extra electron should be distributed into the next subshell of a shell.

            However, the arrangement of electrons becomes complicated for atoms with a high number of electrons.The reason is that once a third-tier subshell is filled up, filling of the fourth and subsequent energy tiers reduces the overall energy compared to the one experienced in the third tier and below(Barrett 97). Also, regardless of the principle’s focus on an electron energy, it can be noted that the angular momentum quantity affects the electron energy, hence, as layers of energy increase, there are cases of overlap on arrangement in the shells, implying that, with reduced energy at higher level, electrons tend to overlap on lower shells or subshells(Burke and Joachain 137). For instance, once the 3d shell is filled, e.g., in the case of Ar., the K atom electrons will be positioned in 4s subshell instead of the 3d subshell as illustrated below;

Figure 1 below is a summary of the order of filling electrons in a shell. The arrows pass through the subshells following their respective filling sequence. As the size of increase, there is some shift in the quantum principle.

Figure 1: Electron filling order in a shell

Energy Diagrams for Electron Configuration

The first part has established how electrons are arranged inside an atom. It is noted that the electrons are orbit inside the shells through a regular pattern. However, this section seeks to understand the reason behind the electrons sticking to such a pattern. According to Burke and Joachain (67-83),   three important principles are followed to ensure that the orbital energy diagram for an electron is appropriately filled. These principles include the Hund’s rule, Pauli Exclusion, and Aufbau principles.  Aufbau principle yields the overall pattern embraced in the order of shell filling. According to the principle (which in German means “building up”), the lowest energy orbitals are filled first (74). When the quantum number increases, the levels of orbital energy increases accordingly because the electron density becomes dispersed far from the nucleus (97). With the exception of hydrogen, there is variation in the energy levels of an electron shell due to constant repulsion between and among the electrons (106).

Consequently, there is an increase in energy level as the quantum number angular momentum, l increases, particularly for orbitals that are sharing similar quantum number principle, n. Figure 2 below demonstrates the Aufbau principle. In the chart, the orbital is represented by a line while a set of lines represents orbital subshells.

Figure 2: Orbital energy configuration for the multi-electron atom

In attribution to the Pauli Exclusion Principle, only a maximum of two electrons should be filled run in each orbital. Hence, it becomes complicated in the case of multiple orbitals of the same energy in a subshell. The same energy level orbitals are called degenerate orbitals. They can be filled using the Hund’s rule, which demands that a single electron be placed first in every degenerate orbital before pairing them in a similar orbital. For example, taking a case of Baron, that contain five electrons, filling begins from the lowest level of energy, that is, 1s orbital where two electrons with opposite spin are placed. Then, it is followed by 2s, orbital that is filled with two electrons, and the last electron is positioned in the previous degenerate orbital as shown in figure 3 below.

Figure 3: Energy configuration for Baron Electrons

Figure 3 is the energy configuration for Baron Atom. It displays how the five electrons are located inside the atom. According to the periodic table element arrangement order, Hund’s rule is applied to include extra electron on the 2p orbital in the increasing number of electrons for thefollowing elements(Barrett 118). As a result, using the three principles makes it easy to study and understand the nature, organization, and positioning of electrons inside an atom.  Also, the three principles enable one to solve the increasing complexity in the organization of electrons inside shells and subshells as the size of the atom increase across the periodic table. Key to understanding their distribution or the arrangement is the levels of energy that determines the quantity of electrons held at each level. The knowledge extends further to understanding the behavior of an atom when subjected to varied sets of temperature, among other physical characteristics that may influence the behavior of an atom.

Works Cited

Barrett, Jack. Atomic structure and periodicity. Hoboken, N.J. Cambridge, UK: Wiley-Interscience Royal Society of Chemistry, 2002. Print.

Burke, P. G., and C. J. Joachain. Theory of electron-atom collisions. New York: Plenum Press, 1995. Print.

Krane, Kenneth S. Modern physics. Hoboken, New Jersey: John Wiley & Sons, Inc, 2020. Print.

Manini, Nicola. Introduction to the physics of matter: basic atomic, molecular, and solid-state physics. Cham: Springer, 2014. Print.