The behavior of electrons is elaborated by other principles of atomic physics such as Hund`s rule and Pauli`s exclusion principle. Hund`s rule states that when several orbitals of the same energy are available, the electrons occupy different orbitals individually before being doubly occupied. When a double assignment occurs, the Pauli exclusion principle requires that electrons occupying the same orbital have different spins (+1/2 and -1/2). Half-filled subshells have weaker electron-electron repulsions in the orbitals, which increases stability. Similarly, fully filled sublayers increase the stability of the atom. Therefore, the electronic configurations of some atoms do not obey the construction principle (depending on the energy difference between the orbitals). Aufbau is a German word meaning “to build” and, unlike many other principles of chemistry, is not the name of a scientist. This principle essentially concerns the filling of electrons in an orbital when writing an electronic configuration. The construction principle states that electrons in the ground state of an atom or ion fill atomic orbitals with the lowest available energy levels before occupying higher levels. For example, subshell 1s is populated before subshell 2s is occupied.
In this way, the electrons of an atom or ion form the most stable electronic configuration possible. An example is the 1s2 2s2 2p6 3s2 3p3 configuration for the phosphorus atom, which means that the 1s subshell has 2 electrons and so on. The construction principle contains some important rules for filling the orbitals in an atom. An atom in its ground state has the lowest energy and is the most stable. The filling of orbitals in the ground state of the atom is carried out according to the construction principle, which is also based on the Pauli exclusion principle and finances the maximum multiplicity and relative energies of the orbitals. The principle states that electrons are gradually added to the different orbitals in ascending order of energy. Electrons first enter the lowest energy orbital available, and then into higher energy orbitals only after the orbitals are filled with lower energy. The electronic configuration of chromium is [Ar]3d54s1 and not [Ar]3d44s2 (as suggested by the design principle). This exception is attributed to several factors, such as the increased stability due to the half-filled subhulls and the relatively small energy difference between the 3d and 4s subhulls. Let us now study the concept of orbitals. In the quantum mechanics model, the specific regions of space around the nucleus where the probability of finding the nucleus is greatest. These regions are expressed by mathematical expressions and are called orbital or more commonly orbital wave functions.
An atom has a large number of orbitals. These orbitals differ quantitatively in shape, size and orientation. A smaller orbital means that the probability of finding the electron near the nucleus is greater. Therefore, Aufbauba`s principle contains some important rules for filling the orbitals in an atom. The following diagram clearly shows how the orbitals are energetically arranged in ascending order. As with most rules, there are exceptions. The half-filled and fully filled subshells d and f give the atoms stability, so the elements of the d and f blocks do not always follow the principle. For example, the expected build configuration for Cr is 4s23d4, but the observed configuration is actually 4s13d5. This actually reduces the electron-electron repulsion in the atom, as each electron has its own seat in the subshell. “Aufbauprinzip” is a German name; It means “construction principle”. There are two other important rules studied by students to fill the orbitals in an atom are the Pauli exclusion principle and Hund`s rule of maximum multiplicity. Not all elements fall into the category of the construction principle.
For example, ruthenium, rhodium, silver and platinum are exceptions to the principle of construction because of the filled or half-filled subhulls. For example, copper is another exception to this principle with an electronic configuration corresponding to [Ar]3d104s1. This is due to the stability of a fully filled 3D underlayment. The building principle means, in simple terms, that electrons are added to orbitals as protons are added to an atom. The term comes from the German word “aufbau”, which means “built” or “building”. The lower electron orbitals fill up before the upper orbitals and “build” the electron shell. The end result is that the atom, ion or molecule forms the most stable electronic configuration. Another content, which was studied by 11 std students, concerned the principle of construction. The construction principle describes the rules according to which electrons are organized in layers and subshells around the atomic nucleus. One electron is spin-up and the other spin-down according to the Pauli exclusion principle. This is called spin matching. It is important to note that each orbital can contain a maximum of two electrons (according to the Pauli exclusion principle).
In addition, the way electrons are filled in orbitals in a single subshell must follow Hund`s rule, i.e. each orbital in a given subshell must be individually occupied by electrons before any pair of electrons in an orbital. The building principle can be used to understand the location of electrons in an atom and their corresponding energy levels. For example, carbon has 6 electrons and its electronic configuration is 1s22s22p2. This diagram is also known as the construction principle diagram and is used to remember the order in which the orbitals are filled. Figure 1. Order of filling of orbitals according to the principle of construction Sulfur has an atomic number of 16, that is, it has 16 electrons in an atom. As mentioned above, the first 2 electrons are occupied by orbital 1s. The next 2 are occupied by the 2s orbitals. The next 6 electrons are occupied by the 2p orbitals. The next 2 electrons are occupied by the 3s orbital and the rest of the last 4 electrons are occupied by the 3p orbitals. Of the 16 electrons, a total of 10 electrons are in the 1st and 2nd shells, i.e.
n = 1 and n = 2 and the last 6 electrons in the 3rd shell, i.e. n = 3. Thus, the valence shell is the 3rd layer and the total number of valence electrons is 6 (2 electrons in 3s and 4 electrons in 3p) in sulfur. We filled electrons according to the construction principle and used Figure 1. The construction principle states that the arrangement of electrons in an atom – the electronic configuration – is best understood when it is constructed from scratch. Although most electronic configurations follow the above order according to the construction principle, there are some exceptions. A handful of elements with atomic numbers greater than 20, such as Cu (copper, atomic number = 29), Cr (chromium, atomic number = 24), Mo (molybdenum, atomic number = 42), etc., are exceptions. These exceptions are due to fully filled or semi-filled atomic orbitals, which are more stable than all partially filled atomic orbitals due to symmetry and exchange energy release. The principle takes its name from the German construction principle, “Aufbauprinzip”, and is not named after a scientist.
It was formulated in the early 1920s by Niels Bohr and Wolfgang Pauli. It was an early application of quantum mechanics to the properties of electrons and explained chemical properties in physical terms. Each electron added is subject to the electric field generated by the positive charge of the atomic nucleus and the negative charge of the other electrons attached to the nucleus. Although in hydrogen there is no difference in energy between orbitals with the same principal quantum number n, this is not true for the outer electrons of other atoms. The construction principle is examined by 11 students according to the rules of placement of orbitals in an atom. This topic is discussed in the chapter on the atomic structure of physical chemistry. This chapter uses the student`s study: the building principle to predict the electronic configurations of atoms and, therefore, explain the arrangement of the periodic table and how electrons are arranged from low to high energy levels.