Classification of Elements: Our surroundings are made up of basic units known as elements. Only 31 elements were known around 1800. Later, with technological advancements, 63 more elements were discovered. As the number of elements grew, it became necessary to classify them on a regular basis.

There are currently 118 elements, 93 of which are found naturally, and the rest were created in laboratories. It was difficult to study these elements independently without first establishing a relationship between them based on their properties. As a result, these elements are organised or classified into different categories or groups based on similarities and differences in their properties, as well as periodicity. Because the elements in the classification of elements are arranged in groups known as periods, this classification system is known as the periodic classification of elements.

What is the Classification of Elements?

The grouping of elements with similar characteristics is known as the Classification of elements. Even though each element is distinct from the others, some elements share similarities. Based on these similarities, scientists were eventually successful in grouping the various elements into groups or chemical families, so that, similar elements were grouped together and dissimilar elements were separated from one another after showing a group. Thus, the Classification of elements leads to the formation of the periodic table.

Need for Classification of Elements

Elements are the basic units of all matter. Some of the advantages of classification of elements are given below:

1. Classification assists us in understanding the properties of elements and their compounds.
2. It is simpler to study, compare, and distinguish the properties of elements and compounds from different groups.
3. The properties of elements and their compounds can be predicted based on their position in the periodic table.

Periodic Classification of Elements Based on Electronic Configuration

The long form of the periodic table is divided into four blocks known as the s, p, d and f-blocks. 

s-block Elements

The elements in the s-block of the periodic table have the last electron filled in valence s-sub-shell of the outermost energy shell. Because the s-subshell can only have two electrons, this block comprises only two groups \(\left({1,2} \right)\). The elements included in group \(1\) have \({\text{n}}{{\text{s}}^{\text{1}}}\) electronic configuration and are called alkali metals. Similarly, group \(2\) contains alkaline earth metals with \({\text{n}}{{\text{s}}^{\text{2}}}\) electronic configuration. 

Thus, the elements of group \(1\) and \(2\), including hydrogen and helium in which the s-orbitals are being progressively filled in are called s-block elements.

General electronic configuration of s-block elements is : \({\text{n}}{{\text{s}}^{\text{1-2}}}\) where \({\text{n}} = 2 – 7\)

p-block Elements

P-block elements are those in which the last electron enters one of the three p-orbitals of their corresponding outermost shells. As the p-subshell can have a maximum of six electrons, there are six groups in this block (\(13\) to \(18\)).

General electronic configuration of p- block elements is : \({\text{n}}{{\text{s}}^{\text{2}}}{\text{n}}{{\text{p}}^{{\text{1-6}}}}\) where \({\text{n}} = 2 – 7\)

The elements belong to s and p-blocks in the periodic table are collectively called Representative elements.

d-block Elements

Elements in which the last electron enters any one of the five d-orbitals of their respective penultimate shells are called d-block elements. Because the d-subshell can have a maximum of five orbitals and ten electrons, the d-block is divided into ten vertical columns or groups numbered \(3\) to \(12\). The d-block elements are known as the transition elements because they have incompletely filled d-orbitals in their ground state.

General electronic configuration of p-block elements is: \(\left( {{\rm{n – 1}}} \right){{\rm{d}}^{{\rm{1 – 10}}}}{\rm{\;n}}{{\rm{s}}^{{\rm{0 – 2}}}}{\rm{\;}}\) where \({\text{n}} = 4 – 7\)

f-block Elements

Elements in which the last electron enters any one of the seven f-orbitals of their respective ante-penultimate shells are called d-block elements. All the f-block elements are also called two such series, each with fourteen elements; in the first series, the filling of electrons takes place in the \({\text{4f}}\)-subshell, known as lanthanoid series (Atomic number \(58-71\)). In the second series, the filling takes place in the \({\text{5f}}\)-subshell, and it is called the actinide series.

General electronic configuration of p- block elements is: \({\text{(n-2)}}{{\text{f}}^{{\text{0-14}}}}{\text{(n-1)}}{{\text{d}}^{{\text{0-2}}}}{\text{n}}{{\text{s}}^{\text{2}}}\) where \({\text{n}} = 6 – 7\)

Periodic Classification of Elements in Periodic Table

Based on nature and characteristics, elements are classified into three main types:

(i) Metals: These comprises \({\text{78% }}\) of all the known elements and appear on the left-hand side of the periodic table. All the s, d, and f block elements are metals. Usually, Metals are solid at room temperature with high melting and boiling points. Except for mercury because it is liquid at room temperature. They are malleable and good conductors of heat and electricity.

(ii) Non-metals: These elements lie on the top right-hand side of the periodic table. The number of non-metal is very few. Non-metals are either solids or gases at room temperature except bromine which is liquid with low boiling and melting point. They are bad conductors of heat and electricity, and they are brittle solids.

(iii) Metalloids: These have specific characteristics common to both metals and non-metals. These are also known as semi-metals.

The elements in the four different blocks of the periodic table can also be classified into four types. They are:

1. Noble gases: These are the elements found in Group \(18\). The general electronic configuration of the elements is \({\text{n}}{{\text{s}}^{\text{2}}}{\text{n}}{{\text{p}}^{{\text{1-6}}}}\) which means that they have completely filled valence s-and p-orbitals (helium with \({\text{1}}{{\text{s}}^2}\) configuration is an exception).

2. Representative elements: These contain both s and p-block elements, except for noble gases, which are classified in their own category.

3. Transition elements:  The transition elements are found in the middle of the periodic table, between the s- and p-block elements, and include all d—block elements. These were separated into \(10 (3–12)\) groups.

4. Inner transition elements: Inner transition elements belong to the f-block of the periodic table and form transition series within the transition elements of d-block.

Classification of Elements and Periodicity in Properties

Periodic properties are properties of elements directly or indirectly related to the electronic configuration of their atoms and show gradation as they move down a group or along a period.

1. Atomic and Ionic Radii

The distance from the centre of the nucleus to the outermost shell containing the electrons is known as atomic radius.

Depending on the nature of the bonding in the atoms, atomic radii is classified into three forms.

Variation of atomic radii in the periodic table
(a) Variation along a period: The atomic radii of elements decreases with an increase in atomic number as we move from left to right in a period. For example, consider the atomic radii of the elements of the second period.

Atom (Period \(2\))	\({\text{Li}}\)	\({\text{Be}}\)	\({\text{B}}\)	\({\text{C}}\)	\({\text{N}}\)	\({\text{O}}\)	\({\text{F}}\)	\({\text{Ne}}\)
Atomic radius	\(152\)	\(111\)	\(88\)	\(77\)	\(70\)	\(74\)	\(72\)	\(160\)

(b) Variation within a group: The atomic radii of elements increases with an increase in atomic number as we move from top to bottom. For example, consider the atomic radii of the members of the alkali metal group.

Alkali metals(Group \(1\))	Atomic radius
\({\text{Li}}\)	\(152\)
\({\text{Na}}\)	\(186\)
\({\text{K}}\)	\(231\)
\({\text{Rb}}\)	\(244\)
\({\text{Cs}}\)	\(262\)

Ionic radii: The effective distance from the nucleus of the ion to which it exerts influence on its electronic cloud is known as the ionic radius.

Variation of Ionic radii in the periodic table

a. Variation of ionic (cationic) radii within a group: Ionic radii of cations increases as we move from top to bottom within a group primarily due to an increase in the number of shells.
b. Variation of ionic (anionic) radii within a group: Ionic radii of anion increase as we move from top to bottom within a group primarily due to an increase in the number of shells.

2. Ionisation Enthalpy

It is the amount of energy required to remove the most loosely bonded electron from an isolated gaseous atom in order to transform it into a gaseous cation (or energy). It is denoted by \({{\rm{\Delta }}_{\rm{i}}}{\rm{H}}{\rm{.}}\)

The following factors determine the ionisation enthalpy:  

1. Atomic size: The ionization energy decreases as the atomic size increases because the electrons are away from the nucleus and thus easy to be removed. It, therefore, increases when the atomic size decreases. 
2. Effective nuclear charge: Ionization energy increases with increased effective nuclear charge because a heavier nuclear charge pulls electrons strongly towards themselves. 
3. Screening effect of the inner electrons: This is also called the shielding effect of the inner shell electrons on the nuclear pull for outer electrons. This effect causes a decrease in the ionization energy because outer electrons are less strongly held.

Variation of Ionic enthalpy in the periodic table
a. Variation along a period: As we move from left to right in a period, the ionization enthalpy increases with an increase in atomic number. This is because the size of the atom decreases with an increase in atomic number, and thus the electrons of the valence shell are closer to the nucleus and are pulled strongly by the protons. Thus, to remove an electron from an atom, more energy is needed. 
b. Variation within a group: The ionization enthalpies keep on regularly decreasing as we move down a group from one element to the other. This is because there is an increase in the number of the main energy shells moving from one element to the other element and there is also an increase in the magnitude of the screening effect due to the gradual increase in the number of inner electrons.

3. Electron Gain Enthalpy

The change in enthalpy occurs when a gaseous atom gains an extra electron to form a monovalent anion, also in the gaseous state, known as electron gain enthalpy. It is denoted by \({{\rm{\Delta }}_{{\rm{eg}}}}{\rm{H}}{\rm{.}}\) 

Variation of Electron Gain enthalpy in the periodic table
a. Variation along a period: The atomic size decreases, and the nuclear charge increases as we walk across a period from left to right. Both of these factors increase the nucleus’s attraction to the incoming electron, and as a result, the electron gain enthalpy decreases from left to right. 
b. Variation within a group: The atomic size and nuclear charge increase as we move down a group. However, increasing atomic size is far more pronounced than the effect of increasing nuclear charge. As atomic size increases, the nucleus’s attraction to the incoming electron lessens, and the electron gain enthalpy becomes less negative.

4. Electronegativity

The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is called its electronegativity.

Electronegativity values also depend on the following factors:
I. Atomic size: As the atomic size increases, electronegativity decreases. This is because the electrons are far from the nucleus, and there is a lesser force of attraction.
II. Nuclear charge: As the nuclear charge increases, electronegativity increases. This is because an increase in nuclear charge causes electrons attraction with greater force.

Variation of Electronegativity in the periodic table
a. Down the group: The electronegativity decreases as the atomic number increases from top to bottom. This is because An atom with a small atomic size has more attraction for the electrons than an atom having a larger atomic size. Since the atomic size increases along with a group, the electronegativity decreases as the atomic number increases from top to bottom.
b. Along a period: Electronegativity increases along a period due to the increase in the nuclear charge. This is because an atom with a small atomic size has more attraction for the electrons than the atom having a larger atomic size (in a chemical bond). Now, across a period, the atomic size decreases. Therefore, electronegativity (attraction for electrons) increases as the atomic number increases from left to right.

Periodic Classification of Elements Class 10 Notes

A periodic table is an arrangement of elements having similar properties placed together. 

Early attempts to classify elements are as follows:

1. Dobereiner’s triads: Dobereiner classified the elements into groups of three elements having similarly in physical and chemical properties.
2. Newland’s Law of Octave: Newland observed that when elements are arranged in increasing order of their atomic mass, every eighth element beginning from any element resembles the first element in its physical and chemical properties.
3. Mendeleev’s Periodic table: Mendeleev arranged the elements  known at the time based on the periodic law;
a) Atomic weight is the fundamental property of an element.
b) The physical and chemical properties of the elements are periodic functions of their atomic weight.
4. Modern periodic table: Moseley arranged elements in the order of increasing atomic number. It consists of seven periods and eighteen groups.

Periodic Classification – Summary

Scientists were eventually successful in grouping the various elements into groups or chemical families. Dissimilar elements were separated from one another after showing a group. As a result, it is simpler to study, compare, and distinguish the properties of elements and compounds from different groups.

FAQs Related to Classification of Elements

Q.1. What are the advantages of the classification of elements?
Ans: It enables scientists to predict the properties of elements and their compounds based on their locations in the periodic table, and vice versa. It becomes easier to study, understand, compare, and contrast the relative properties of elements and their compounds from various classes.

Q.2. What is the use of the classification of elements?
Ans: Uses of Periodic classification are;
a. Classification assists us in understanding the properties of elements and their compounds.
b. It is simpler to study, compare, and distinguish the properties of elements and compounds from different groups.
c. The properties of elements and their compounds can be predicted based on their position in the periodic table.

Q.3. What are the \(4\) types of elements?
Ans: s, p, d and f block elements are the four types of elements.

Q.4. What is the classification of elements on the basis of properties?
Ans: Based on the properties, elements are classified into \(3\) types. They are metals, non-metals and metalloids.

Q.5. Why do we classify elements?
Ans: Scientists gained more knowledge on the properties of new elements as they were found. It was discovered that organising all of the information or properties of these elements was difficult. As a result, scientists began to look for patterns in their properties in order to classify all of the known elements and make their research easier.