The electronic configuration, Also called electronic structure, it is the arrangement of electrons in energy levels around an atomic nucleus. According to the old Bohr atomic model, electrons occupy various levels in orbits around the nucleus, from the first shell closest to the nucleus, K, to the seventh shell, Q, which is the furthest from the nucleus..
In terms of a more refined quantum mechanical model, the K-Q shells are subdivided into a set of orbitals, each of which can be occupied by no more than one pair of electrons..
Commonly, the electron configuration is used to describe the orbitals of an atom in its ground state, but it can also be used to represent an atom that has ionized into a cation or anion, compensating for the loss or gain of electrons in their respective orbitals..
Many of the physical and chemical properties of elements can be correlated with their unique electronic configurations. Valence electrons, the electrons in the outermost shell, are the determining factor for the unique chemistry of the element..
Before assigning the electrons of an atom to orbitals, one must familiarize oneself with the basics of electron configurations. Each element in the Periodic Table consists of atoms, which are made up of protons, neutrons, and electrons..
Electrons exhibit a negative charge and are found around the nucleus of the atom in the orbitals of the electron, defined as the volume of space in which the electron can be found within 95% probability..
The four different types of orbitals (s, p, d, and f) have different shapes, and one orbital can hold a maximum of two electrons. The p, d, and f orbitals have different sublevels, so they can hold more electrons.
As indicated, the electron configuration of each element is unique to its position in the periodic table. The energy level is determined by the period and the number of electrons is given by the atomic number of the element.
Orbitals at different energy levels are similar to each other, but occupy different areas in space..
The 1s orbital and the 2s orbital have the characteristics of an s orbital (radial nodes, spherical volume probabilities, they can only contain two electrons, etc.). But, because they are at different energy levels, they occupy different spaces around the nucleus. Each orbital can be represented by specific blocks on the periodic table..
Block s is the region of alkali metals including helium (Groups 1 and 2), block d is the transition metals (Groups 3 to 12), block p is the elements of the main group of Groups 13 to 18 , And the f block are the lanthanide and actinide series.
Aufbau comes from the German word "Aufbauen" which means "to build". In essence, by writing electron configurations we are building electron orbitals as we move from one atom to another..
As we write the electron configuration of an atom, we will fill in the orbitals in increasing order of atomic number.
The Aufbau principle originates from the Pauli exclusion principle which says that there are no two fermions (eg electrons) in an atom. They can have the same set of quantum numbers, so they have to "stack" at higher energy levels.
How electrons accumulate is a topic of electron configurations (Aufbau Principle, 2015).
Stable atoms have as many electrons as protons do in the nucleus. Electrons gather around the nucleus in quantum orbitals following four basic rules called the Aufbau principle..
The second and fourth rules are basically the same. An example of rule four would be the 2p and 3s orbitals.
A 2p orbital is n = 2 and l = 2 and a 3s orbital is n = 3 and l = 1. (N + l) = 4 in both cases, but the 2p orbital has the lowest energy or lowest n-value and will fill before the layer 3s.
Fortunately, the Moeller diagram shown in Figure 2 can be used to do electron filling. The graph is read by running the diagonals from 1s.
Figure 2 shows the atomic orbitals and the arrows follow the way forward.
Now that it is known that the order of the orbitals is filled, the only thing left is to memorize the size of each orbital.
S orbitals have 1 possible value of ml to hold 2 electrons
P orbitals have 3 possible values of ml to hold 6 electrons
D orbitals have 5 possible values of ml to hold 10 electrons
F orbitals have 7 possible values of ml to hold 14 electrons
This is all that is needed to determine the electronic configuration of a stable atom of an element..
For example, take the element nitrogen. Nitrogen has seven protons and therefore seven electrons. The first orbital to fill is the 1s orbital.
An s orbital has two electrons, so there are five electrons left. The next orbital is the 2s orbital and contains the next two. The final three electrons will go to the 2p orbital that can hold up to six electrons (Helmenstine, 2017).
Electron configurations play an important role in determining the properties of atoms.
All the atoms of the same group have the same external electronic configuration with the exception of the atomic number n, which is why they have similar chemical properties.
Some of the key factors that influence atomic properties include the size of the largest occupied orbitals, the energy of the higher-energy orbitals, the number of orbital vacancies, and the number of electrons in the higher-energy orbitals..
Most atomic properties can be related to the degree of attraction between the outermost electrons to the nucleus and the number of electrons in the outermost electron shell, the number of valence electrons.
The electrons of the outer shell are the ones that can form covalent chemical bonds, they are the ones that have the ability to ionize to form cations or anions and they are the ones that give the oxidation state to the chemical elements.
They will also determine the atomic radius. As n gets larger, the atomic radius increases. When an atom loses an electron, there will be a contraction of the atomic radius due to the decrease in negative charge around the nucleus..
The electrons of the outer shell are those that are taken into account by the valence bond theory, crystalline field theory and molecular orbital theory to obtain the properties of the molecules and the hybridizations of the bonds..
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