What is the Electronic Configuration of Elements?
Each atom belonging to every single group and period has a characteristic electronic configuration. This electronic configuration of the elements determines their several properties. It plays a major role in determining the chemical properties of the atoms such as valency, conductance, reactivity. If we look at it on a wider scale, electronic configuration determines every single chemical property of an element. Each element, in the ground state, has a unique electronic configuration and hence, each one of them differs from the other in many ways.
Electrons, although they are so tiny, have a detrimental role in several reactions. The electronic configuration also determines which element would combine with which other element and in what ratio. Hence, the electronic configuration does play a huge role in the existence of every object in the surrounding.
Electronic Configuration in Periods of the Periodic Table
The elements in a periodic table are arranged in:
Periods
Groups.
(image will be uploaded soon)
The periods are basically the horizontal rows and the groups are the vertical columns of the modern periodic table. As we see in the image above, the modern periodic table has 7 periods and 18 groups. Across a period the atomic number of every consecutive element differs by 1. The electronic configuration between consecutive elements differs by only 1, across a period.
The value of the principal quantum number, that is n is incremented by 1 as we traverse across a period. As we traverse from left to right across a period, the Electronic Configuration of the elements increases, and hence the energy level of the atoms also increases.
The first period is the smallest period and hosts only two elements, hydrogen and helium. These two elements do not have a P subshell and have only an S subshell.
The electronic configuration of hydrogen is 1s¹ whereas the electronic configuration of helium is 1s².
As we proceed down a group, the period number increases and the number of subshells also increases.
The second period begins with lithium and boron which have 3 and 4 electrons respectively.
The 6th and 7th periods could be classified as the longest periods in the modern periodic table, as they also host lanthanides and actinides.
Electronic Configuration in Groups
As we have seen earlier, the modern periodic table consists of 18 groups. As we traverse down a group the electronic configuration in groups changes hugely. A new subshell is added to the elements as we traverse down a group. All the elements in a group have similar properties and also have the same valency. The elements belonging to a particular group have several similarities in their chemical properties and also form similar types of bonds, however, there are some exceptions.
Elements of each group have their characteristic features. A list of the features is listed below.
Alkali Metals
Group 1: Group 1 consists of alkali metals. The valency of these metals is 1 and these metals are highly reactive. The metals of group 1 include Lithium, Sodium, Potassium, Rubidium, Cesium, and Francium.
Group 2: Group 2 elements are called alkali earth metals. These metals are named so because they are found in the earth's crust. The group 2 elements include Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), and Barium (Ba).
Transition Metals
Transition metals include metals from group 3, group 4, group 5, group 6 until group 12. The transition metals constitute one of the largest sets of the periodic table. The elements classified as transition metals are strong and durable. A few of the prominent metals of the transition metal group include iron, manganese, copper, precious metals such as gold, silver, and platinum. These metals are also called d block elements.
Metalloids
Metalloids are a peculiar set of elements that have properties of both metals and nonmetals. The chemical and physical properties of metalloids are a perfect mixture of both metals and nonmetals. They are distinguished by a characteristic zig-zag line in the right-hand side of the periodic table ranging from group 13 to group 16. The following elements are classified as metalloids.
Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po).
Nonmetals
Nonmetals are the metals on the left-hand side of the periodic table. These elements are called so since they aren't metallic in nature. The prominent non-metals include the group 15 elements which have carbon. Carbon is one of the most important elements. This is because most naturally occuring materials consist of carbon. The non-metals also include nitrogen, oxygen, and other gasses. Life isn't possible without any of these gases. Halogens and noble gasses also come under nonmetals. Halogens are highly reactive gasses whereas noble gasses are the least reactive metals.
History of Electronic Configuration
The construction of the periodic table started during the 1860s when the Russian scientist Dmitry Mendeleyev started making a systematic analysis of the link among the characteristics of elements. He established the periodic law and designed the tabular organization of elements. This technique made it feasible to examine various sorts of chemical connections that before had been explored solely in isolation. It was just in the mid-1900s that headway was achieved in understanding the periodic rule in terms of the electrical configuration of atoms.
Open and Closed Shells
In the setting of atomic orbitals, an open shell is a virtual shell that is not entirely occupied with electrons or that has not supplied all of its valence electrons via chemical interactions with other atoms or molecules within a chemical process. Alternatively, a closed shell is formed with a totally filled valence shell. This arrangement is quite stable. For molecules, "open shell" indicates that there are unattached electrons. In molecular orbital theory, this results in molecular orbitals that are individually filled. In computational chemistry applications of molecular orbital theory, open-shell molecules have to be treated using either the limited open-shell Hartree–Fock technique or the unconstrained Hartree–Fock method. Open shell molecules are more challenging to examine electronically.
Applications
The most usual application of electron arrangements is in the rationalization of chemical characteristics. This is the case in both inorganic and organic chemistry. In essence, electron configurations, coupled with a reduced version of the molecular orbital theory, have become the contemporary equivalent of the valence notion, defining the quantity and kind of chemical bonds that an atom might be anticipated to form. For atoms or molecules containing more than one electron, the movement of electrons are coupled and such a representation is no longer precise. A very high variety of electronic combinations are required to accurately represent any multi-electron circuit, and no energy can be linked with one single form. However, the electronic wave pattern is frequently controlled by a very small number of combinations. Hence, the idea of the electronic configuration remains crucial for multi-electron systems.
FAQs on Electronic Configuration Periods Groups – Explanation and FAQs
1. What role does the Electronic Configuration of Elements play?
Electronic configuration of elements plays a key role in determining their chemical properties. Although electrons are tiny, they practically determine the nature of the elements. They determine the chemical properties of the element- the valency, the ionization potential, the ionization enthalpy, chemical bonding, and almost every single property. The lack of an electron makes the element an electron acceptor, whereas having an excess electron makes it an electron donor. Hence, electronic configuration is a deciding factor like the electron.
2. What is the Electronic Configuration of the first 30 elements?
Listed below are the electronic configurations of 30 elements, this also includes the electronic configuration of the first 20 elements.
3. What is the periodic table?
The periodic law is a key principle of chemistry that has enabled scientists to explore the characteristics of elements and how they correlate to each other. The rule asserts that chemical elements show a repetition of attributes when they are placed in sequence of increasing atomic numbers. An atomic number is the total number of protons in the atomic nucleus. This sequential arrangement is known as the periodic table.
The columns of the periodic table are termed groups. Elements of the identical group in the list have the same amount of electrons in the outermost shells of their atoms and form bonds of the identical type.
The straight rows are called periods. Periods relate to the connection of orbitals or probable locations in which electrons will be located, inside the outermost shell of the atom. Subsequent periods along the table correlate to atoms having a more electron-rich core of inner shells.
4. What are shells?
An electron shell is the collection of permissible states that have the identical primary quantum number, n (the number preceding the letter in the orbital designation), that electrons may inhabit. An atom's nth electron shell can house 2n2 electrons, for instance, the first shell can accept 2 electrons. The second shell has 8 electrons, and the third shell has 18 electrons, and so on. The factor of two emerges because the allowable states are doubled owing to particle spin atomic orbital permits up to two identically similar electrons with opposing spin, one with a spin +1⁄2 (typically symbolized by an up-arrow) and one with a spin of −1⁄2 (with a down-arrow) (with a down-arrow).
5. What are subshells?
A subshell is the collection of units described by a shared azimuthal quantum number, l, within a shell. The number of l is in the range from 0 to n − 1. The numbers l = 0, 1, 2, 3 equate to the s, p, d, and f labels, correspondingly. Like if a 3d subshell has n equals 3 and l equals 2. The greatest number of electrons that may be inserted in a subshell is determined by 2(2l + 1). The proportions of electrons that can fill each shell and each subshell originate from the formulas of quantum physics in specific. The Pauli Exclusion Principle states that no two electrons can have the same values of the four quantum numbers in the same atom.
