Periodic Classification of Elements

What does Classification of Elements mean?

Arranging all the known elements so that elements with identical characteristics known as periodic properties or trends are placed together in a sequence is known as the classification of elements.

The Need of Classification of Elements

Around the 17th century, when the conceptualization of elements and their discovery had just begun, only a few elements were known. But as time proceeded, plenty of elements started to come into the picture, and it was difficult for scientists to study each element and their respective compounds individually. Hence, the idea of Classification of Elements came up, which enabled the study of a similar group of elements without any hassle.

Genesis of Periodic Classification

Right from the 19th century, several attempts have been made for an apt classification of elements. Here are some of the most noteworthy ones – 

Dobereiner’s Law of Triads: 

Around 1815, a German chemist Johann Wolfgang Döbereiner, gave the Law of Triads according to which elements of similar properties were arranged in groups of 3, such that the atomic mass of 2nd element is an arithmetic mean of the atomic weights of other 2 elements, and the properties of 2nd elements were in between those of end members. But this law was soon rejected as it applied only to a limited number of elements.

De-Chancourtois’ Telluric Helix: 

In 1862, a French geologist, De Chancourtois, attempted to classify elements according to their atomic masses. He took a vertical cylinder of 16 central lines with a helix drawn at 45° around the cylinder and arranged elements around it. It was observed that the elements in a vertical line showed nearly the same properties. However, this classification was not accepted as it was physically inept.

Newland’s Law of Octaves: 

In 1864, a British chemist, John Newlands, presented the law of octaves according to which, “The elements if arranged in increasing atomic mass, the eighth element was the repetition of the first one like 8th note on the musical scale”. However, it was rejected as it didn’t apply to heavier elements and didn’t leave any space for new elements.

 Lothar Meyer’s Curve: 

In 1869, a German chemist, Lothar Meyer, proposed that “The physical properties of the elements are a periodic function of their atomic masses.” He presented this classification in the form of a curve between atomic volumes and atomic masses of elements wherein he made these observations – 

• Alkali metals were positioned at the maxima of the curve.

• Alkaline earth metals occupy positions at about the mid-points on the descending parts of the curve. 

• Halogens occupy positions on the ascending parts of the curve.

• Transition elements occupy the minima of the curve.  

           However, it lacked practical utility as it was difficult to keep the various portions of the curve in mind.

Mendeleev’s Periodic Law: 

In 1869, a Russian chemist Dmitri Mendeleev proposed the periodic law, which stated that ‘The properties of the elements are a periodic function of their atomic masses, emphasising chemical properties.’ Aiding this classification, he created a tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of elements, and called it the ‘Periodic Table’ Some key highlights of this table are as follows – 

• To accommodate 63 known elements of then, 90 spaces were provided, wherein the empty spaces were set for the unknown, to be discovered elements. For example, gallium and germanium were named eka-aluminium and eka-silicon in the table since they weren’t discovered yet. Their properties were predicted by Mendeleev and matched the properties of elements when discovered.

• Elements with identical properties were present in vertical columns called groups. The horizontal rows were called periods.

• Sometimes, properties of elements were emphasised instead of atomic masses. For instance, iodine having a lower atomic mass than tellurium was placed ahead of tellurium because iodine showed similarities with fluorine, chlorine, and bromine

• The table consisted of defects such as anomalous pairs, isotopes’ position, lanthanides and actinides’ position, etc.

Modern Periodic Law and Modified Form of Mendeleev’s Periodic Table: 

In 1913, an English physicist Henry Moseley suggested that the basis of classification of elements should be the atomic numbers of the elements instead of the atomic masses of the elements. He thus modified the periodic law as “The physical and chemical properties of elements are periodic functions of their atomic numbers.” The table thus obtained is called the modern Mendeleev’s Periodic Table. The main characteristics of the modern periodic table are – 

• The table consists of seven horizontal rows called ‘periods’ and nine vertical columns called ‘groups.’ The groups are marked from 0 to VIII. Each group from I to VII is divided into 2 subgroups, A and B.

• Every period starts with a member of the alkali group and ends with a member of the zero group. The first period, however, starts with hydrogen.

• As we move across the periods and groups of the periodic table, we observe gradual changes in a few properties, and these properties are –

•   Valency: if we take oxygen as standard, valency increases from 1 to 8 as we move across a period. Thus, group number indicates the valency of the elements belonging to that group if oxygen is taken as a standard. 

• Metallic nature: As we move across right in a period, the non-metallic nature of the element increases, i.e., metallic nature decreases. 

• Electronegative/ electropositive character: As we move across right in a period, the electronegative nature of an element increases, i.e., electropositive nature decreases. 

• Oxidising/ reducing nature: As we move across the right period, oxidation increases, i.e., reducing nature decreases. 

• Nature of oxides: As we move across right in a period, the acidic nature of oxides increases, i.e., the basic nature of oxides decreases. 

However, due to defects like the incomprehensible position of hydrogen in the periodic table position of lanthanides and actinides, Mendeleev’s modern periodic table was not settled for either.

Extended Or Long Form Of Periodic Table: 

To remove the defects of Mendeleev’s periodic table, several tables have been presented for the classification of elements. The best table out of these is the extended form of the periodic table. This table follows Bohr’s scheme and Aufbau’s principle of arranging elements into four types based on their electronic configurations. 

• s-block: The differentiating electron is present in the ‘ns’ energy shell in these elements. This block lies on the extreme left of the periodic table with Group IA with ns1 configuration known as alkali metals and Group IIA with ns2 configuration known as alkaline earth metals.

• p-block: The differentiating electron is present in the ‘np’ energy shell in these elements. This block lies at the extreme left of the periodic table, and its valence configuration varies from ns2np1 to ns2np6. It contains the groups – IIIA, IVA, VA, VIA, VIIA, 0 (Zero group)

• d-block: The differentiating electron is present in the ‘(n-1)d’ energy shell in these elements. This block lies at the centre of the periodic table, and its valence configuration varies from (n-1)d1ns2 to (n-1)d10ns2. It contains IIIB, IVB, VB, VIB, VIIB, VIII, IB, and lIB. These are known as transition elements.

• f-block: The differentiating electron is present in the ‘(n-2)f’ energy shell in these elements. This block lies at the centre of the periodic table and their valence configuration varies from (n-2)f1(n-1)d1ns2 to (n-2)f14(n-1)d1ns2. All f-block elements belong to the 3rd group – the elements are accommodated in the 3rd group but written separately in two horizontal rows below the periodic table. 

Periodic Properties In Long Form Of Periodic Table: 

• Screening effect or Shielding effect: A decrease in the force of attraction exerted by the nucleus on the valency electrons due to the presence of electrons in the inner shells, is called screening effect or shielding effect. It is observed that the magnitude of screening effect in the case of s and p-block elements increases in a period as well as in a group as the atomic number increases.

• Effective Nuclear Charge is the net amount of positive (nuclear) charge experienced by an electron in an atom. It is observed that as we move towards right in a period, effective nuclear charge increases. It almost remains the same in the s and p block across a group.

• Atomic Radius is the distance between the centre of an atom to its outermost shell. As we move towards the right in a period, the atomic radius decreases in s and p block elements; the atomic radius increases as we move down a group.

• Ionisation Enthalpy: it is the least amount of energy required to remove the outermost electron from an isolated atom in the gaseous state of an element to convert it into a gaseous monovalent positive ion. It decreases with an increase in atomic radius and increases in nuclear charge.

• Electron gain enthalpy is the energy released when a neutral isolated gaseous atom accepts an electron to form an anion. Generally, its value gets less negative in going from top to bottom in a group and more negative in going from left to right in a period.

• Electronegativity is the relative tendency of an atom to attract the shared electron pair towards itself. As we move right, its value increases, while in a group from top to bottom, its value decreases.

Conclusion:

By reading the above article, one can go through the various periodic properties and their trends across the periodic table and understand the importance of classification in such a manner.