Electronic Configurations and Types of Elements

Electronic Configurations and Types of Elements: The electronic configuration of atoms, along with the Aufbau (build-up) principle, provide a theoretical foundation for the periodic classification of elements. A group or family of elements in a vertical column of the Periodic Table has similar chemical behaviour. Since these elements have the same number and distribution of electrons in their outermost orbitals, they are comparable. Let’s explore how elements are classified based on their electronic configurations.

What is Electronic Configuration?

The distribution of electrons in an element’s atomic orbitals is described by its electronic configuration. The electronic configurations of an atom follow a standard nomenclature in which all electron-containing atomic subshells are arranged in a sequence (with the number of electrons they possess written in superscript). The electronic configuration of sodium, for example, is \({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{1}}}\)

On the other hand, the standard notation often yields lengthy electron configurations (especially for elements having a relatively large atomic number). An abbreviated or condensed notation may be used instead of the conventional notation in several cases. The series of completely filled subshells that correspond to a noble gas’s electronic configuration is replaced with the noble gas’s symbol in square brackets in the abbreviated notation. As a result, sodium’s shortened electron configuration is \({\rm{[Ne]3}}{{\rm{s}}^{\rm{1}}}\) (the electron configuration of neon is \({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}},\) which can be abbreviated to \(\left. {{\rm{[He]2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}} \right){\rm{.}}\)

The elements can be classified into four groups based on their electronic configuration: \({\rm{s,p,d,}}\) and \({\rm{f}}{\rm{.}}\) The name of the orbital that receives the valence electron is the basis for this distinction. But there are exceptions too.

Although helium technically belongs to the \({\rm{s – }}\)block, it is placed in the \({\rm{p – }}\)block with other group \(18\) elements because it has a completely filled valence shell \(\left( {{\rm{1}}{{\rm{s}}^{\rm{2}}}} \right)\) and hence shows properties similar to other noble gases.

Hydrogen is the other exception. Because it only has one \({\rm{s – }}\)electron, it belongs to group \(1\) (alkali metals). It can also gain an electron to form a noble gas configuration, making it similar to group \(17\) (halogen family) elements in behaviour. Because it is a special case, we shall place hydrogen separately at the top of the Periodic Table.

s-block Elements

The elements of Group \({\rm{1A}}\) (alkali metals) and Group \({\rm{IIA}}\) (alkaline earth metals) with \({\rm{n}}{{\rm{s}}^{\rm{1}}}\) and \({\rm{n}}{{\rm{s}}^{\rm{2}}}\) outermost electronic configuration constitute the \({\rm{s – }}\)block Elements. These are present in the extreme left side of the Periodic table.

\({\text{S}}\)-block elements have the following characteristics:

• They are soft metals with low melting and boiling temperatures.
• They have a metallic character, and metal reactivity increases down the group.
• They have low ionisation enthalpies and are strongly electropositive metals.
• These metals have valencies of \(+1\) (in the case of alkali metals) and \(+2\) (in the case of other metals) (in the case of alkaline earth metals).
• The flame of most of the metals in this block has a distinct colour.
• These metals are strong reducing agents and excellent heat and electricity conductors.
• They rapidly lose the outermost electron(s) to create a \(1+\) ion (in the case of alkali metals) or a \(2+\) ion (in the case of alkaline earth metals). 

p-block Elements

These elements are present on the right side of the periodic Table and constitute the groups \({\rm{IIIA}}\) to \({\rm{VIIA}}\) and zero groups of the modern Periodic Table. The majority of these elements are metalloids and non-metals; however, there are a few metals among them. The valence shell’s \({\rm{p}}\) orbital is where these atoms’ last electron enters. The valence shell electronic configuration of \({\rm{p – }}\)block elements is \({\rm{n}}{{\rm{s}}^{\rm{2}}}{\rm{n}}{{\rm{p}}^{{\rm{1 – 6}}}}\) where \({\text{n}} = 2\) to \(7\)

Characteristics of \({\text{p}}\) − block elements are as follows:

• It constitutes both metals & non-metals. 
• They typically form covalent compounds, and their metallic character diminishes from left to right along with the period and increases from top to bottom within a group.
• When compared to \({\rm{s – }}\)block elements, the ionisation energy is higher.
• In a group, reducing characters increase from top to bottom, whereas oxidising characters increase from left to right in a period.
• At the end of each period, a noble gas element with a closed valence shell \({\rm{n}}{{\rm{s}}^{\rm{2}}}{\rm{n}}{{\rm{p}}^{\rm{6}}}\) configuration.
• Together with the \({\rm{s – }}\)Block Elements, these elements are known as Representative Elements or Main Group Elements.
• All of the orbitals in the valence shell of noble gases are totally occupied with electrons, and altering this stable configuration by adding or removing electrons is extremely difficult. As a result, the noble gases have an extremely low chemical reactivity.
• Two chemically significant non-metal groups come before the noble gas family. They are chalcogens (Group \(16\)) and halogens (Group \(17\)). They can easily add one or two electrons to achieve a stable noble gas configuration. As we move from left to right across a period, the non-metallic character increases, while the metallic character increases as we move down the group.

d-block Elements (Transition Elements)

These elements are found in Groups \({\rm{IIIB}}\) to \({\rm{VIIB,VIII,IB,}}\) and \({\rm{IIB}}\) of the modern Periodic Table. These elements are positioned at the centre between the \({\rm{s}}\) and \({\rm{p}}\) block of the Periodic Table. The filling of inner \({\text{d}}\) orbitals by electrons distinguishes these elements, which are referred to as \({\rm{d – }}\)Block Elements. The general outer electronic configuration of these elements is \({\rm{(n – 1)}}{{\rm{d}}^{{\rm{1 – 10}}}}{\rm{n}}{{\rm{s}}^{{\rm{1 – 2}}}}\) where \({\rm{(n = 4}}\) to \(7.\) There are four series of \({\rm{d – }}\)block elements, which are \({\text{3d}}\) series \( – {\mathop{\rm Sc}\nolimits} (21)\) to \({\rm{Zn}}(30),4\;{\rm{d}}\) series \({\rm{ – Y(39)}}\) to \({\rm{Cd}}(48)\) & \({\rm{5d}}\) series \({\rm{ – La}}(57),{\rm{Hf}}(72)\) to \({\rm{Hg}}(80),6\;{\rm{d}}\) series \({\mathop{\rm Ac}\nolimits} (89),{\mathop{\rm Rf}\nolimits} (104) \ldots \ldots \ldots \ldots \ldots \ldots ..{\mathop{\rm Cn}\nolimits} (112)\)

Elements of the \({\rm{d – }}\)block have the following characteristics:

• These are hard, ductile, and malleable metals, with high melting and boiling points.
• The ionisation energy is between \({\rm{s}}\) and \({\rm{p – }}\)block elements.
• They have a wide range of oxidation states and are good heat and electricity conductors.
• They can form both ionic and covalent compounds, which are usually coloured and paramagnetic.
•  The majority of transition metals combine to form alloys.

Most of the features of transition elements are absent in \({\rm{Zn}},{\rm{Cd}},\) and \({\rm{Hg,}}\) which have the electronic configuration \({\rm{(n – 1)}}{{\rm{d}}^{{\rm{10}}}}{\rm{n}}{{\rm{s}}^{\rm{2}}}{\rm{.}}\)

Transition metals, in a sense, serve as a link between the chemically active metals of the \({\rm{s – }}\)block elements and the less active elements of Groups \(13\) and \(14,\) giving them the name “Transition Elements.”

f-block Elements (Inner-Transition Elements)

These are located beneath the main periodic Table and are primarily connected to group \({\rm{IIIB,}}\) or the periodic Table’s third group. The \({\rm{4f}}\) series – Lanthanides (\(14\) elements \({\rm{Ce}}(58)\) to \({\mathop{\rm Lu}\nolimits} (71)\) and the \({\rm{5f}}\) series – Actinides \((14\) elements \({\rm{Th(90)}}\) to \({\mathop{\rm Lr}\nolimits} (103))\) are the two series of \({\rm{f – }}\)block elements \((103).\) The general outer electronic configuration of \({\rm{f}}\) block elements is \({\rm{(n – 2)}}{{\rm{f}}^{{\rm{0 – 14}}}}{\rm{(n – 1)}}{{\rm{d}}^{{\rm{0 – 1}}}}{\rm{n}}{{\rm{s}}^{\rm{2}}}.\)

\({\rm{f – }}\)block elements have the following characteristics:

• They are heavy metals with high melting and boiling temperatures; 
• Their compounds have a wide range of oxidation states; and 
• The majority of the actinide elements are radioactive.

The \({\rm{f – }}\)orbital is filled with the last electron added to each atom. As a result, these two sets of elements are known as the Inner-Transition Elements (\({\rm{f – }}\)Block Elements). They’re all made of metal. The characteristics of the elements in each series are relatively similar.

Metals, Non-metals and Metalloids

Besides classifying elements into \({\rm{s – , p – , d – ,}}\) and (\({\rm{f – }}\)blocks, they can also be classified into Metals and Non-metals based on their properties.

Metals are the elements that appear on the left side of the Periodic Table. Some of their properties are-

• These are usually present in the solid form at room temperature with mercury as an exception.
• Metals have high melting and boiling points in general, with the exception of gallium and caesium, which have very low melting points (\({\rm{303K}}\) and \({\rm{302K,}}\) respectively).
• They are excellent heat and electricity conductors.
• They have the properties of malleability (they may be hammered into thin sheets) and ductility (can be drawn into wires).

Non-metals are located at the top right-hand side of the Periodic Table. Some of its properties are-

• As one moves along the horizontal rows in the Periodic Table, the property of elements change from metallic on the left to non-metallic on the right. These are usually solids or gases at room temperature with low melting and boiling points, except boron and carbon.
• They are poor conductors of heat and electricity. 
• These solids are brittle and are neither malleable nor ductile. The non-metallic character decreases moving down the group; as one goes from left to right across the Periodic Table.

Moving along the rows, the change from the metallic property of elements changes gradually to its non-metallic character, and there are certain elements that exhibit the characteristic of both metals and non-metals. These elements are silicon, germanium, arsenic, antimony and tellurium and are represented in a zigzag pattern. These elements run diagonally across the Periodic Table and are called Semi-metals or Metalloids.

Prediction of Group, Period and Block of a Given Element

The electronic configuration of an element helps us to predict its group, period & block by the following ways-

• The period of an element is given by the principal quantum number \(\left( {\rm{n}} \right)\) of the valence shell.
• An element’s block is indicated by the orbital containing the valence electron.
• The number of electrons in the valence shell or a penultimate shell \(\left( {{\rm{n – 1}}} \right)\) can be used to predict an element’s group: 

(a) The number of valence electrons is equal to the group number for \({\rm{s – }}\)block elements.
(b) For \({\rm{p – }}\)block elements, group number equals \(10+\) valence shell electron count.
(c) For \({\rm{d – }}\)block elements, the number of electrons in the \({\rm{(n – 1)d – }}\)subshell \(+\) the number of electrons in the valence shell equals the group number (\({\text{n}}\)th shell). Alternatively, group number \(=\) number of electrons in (penultimate shell \(+\) valence shell) \(-8\) 
(d) The group number for \({\text{f}}\) block elements is always \({\rm{IIIrd/3B}}{\rm{.}}\)

Summary

In the modern Periodic Table, elements are arranged according to their increasing atomic numbers. The atomic number, also known as the proton number, forms the basis for categorising elements into groups and periods in the Periodic Table. The classification is done according to the electronic configuration of different elements. In this article, we learnt the how and why elements are classified into metals, non-metals and metalloids, their characteristic properties, and how to predict the group, period and block of a given element. We also learnt about different blocks in the Periodic Table and the electronic configuration of their outermost shells.

FAQs

Q.1. How do you write electronic configuration?
Ans: While writing the electronic configuration, we write the energy level (the period) first, then the subshell to be filled, and finally the superscript, which is the number of electrons in that subshell. The atomic number \({\rm{Z}}\) represents the total amount of electrons.

Q.2. What is the electronic configuration of copper?
Ans: The electronic configuration of copper is \({\rm{1}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{6}}}{\rm{3}}{{\rm{d}}^{{\rm{10}}}}{\rm{4}}{{\rm{s}}^{\rm{1}}}\).

Q.3. Why are electronic configurations important?
Ans: The atom’s valence electrons provide an insight into the chemical behaviour of elements. It also aids in the classification of elements into distinct blocks (such as the \({\rm{s – }}\)block elements, the \({\rm{p – }}\)block elements, the \({\rm{d – }}\)block elements, and the \({\rm{f – }}\)block elements).

Q.4. What are the three rules that must be followed while writing the electronic configuration of elements?
Ans: We follow the three important rules: Aufbau Principle, Pauli-exclusion Principle, and Hund’s Rule.

Q.5. What is meant by the electronic configuration of an element?
Ans: An atom’s electron configuration is a representation of how electrons are distributed among the orbital shells and subshells. Many of an element’s physical and chemical features can be linked to its unique electronic configuration.

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