The 19th century was marked by the discovery of numerous elements, necessitating a systematic method for their classification. Studying each element individually was becoming a challenging task.
Various attempts were made to categorize these elements, with notable ones being ‘Dobereiner’s Triads’ and ‘Newland’s Octaves’.
Johann Wolfgang Dobereiner, a German chemist, proposed the classification of elements into groups of three, known as ‘triads’. He suggested that in these triads, the atomic mass of the middle element would be approximately equal to the mean of the atomic masses of the other two elements.
For instance, consider a triad consisting of lithium, sodium, and potassium. Lithium has an atomic mass of 6.94 and potassium has an atomic mass of 39.10. The middle element, sodium, has an atomic mass of 22.99, which is nearly equal to the average of the atomic masses of lithium and potassium (23.02).
However, Dobereiner’s Triads had its limitations:
In 1866, John Newlands, an English scientist, arranged the 56 known elements in the order of increasing atomic mass. He noticed a pattern where every eighth element had similar properties to the first. This pattern is illustrated below.
This observation led to the formulation of Newland’s Law of Octaves. According to this law, when elements are arranged in the order of increasing atomic mass, two elements separated by seven others exhibit similar properties.
Newland’s Octaves, however, had its limitations:
In 1869, Dmitri Ivanovich Mendeleev, a Russian chemist, introduced his periodic table. He observed that the physical and chemical properties of elements were periodically related to their atomic masses.
According to the Periodic Law, also known as Mendeleev’s Law, the chemical properties of elements are a periodic function of their atomic weights.
Mendeleev’s Periodic table had its advantages:
However, Mendeleev’s Periodic table had its limitations:
These classification methods laid the foundation for the modern periodic table. Dmitri Mendeleev, who is often referred to as the Father of the Modern Periodic Table, made the most significant contribution. The modern periodic law is also named after him in his honor.
In 1913, English physicist Henry Moseley studied the wavelength of the characteristic x-rays using different metals as anti-cathode. He found that the square root of the frequency of the line is related to the atomic number. Based on these observations, Moseley proposed the modern periodic law, which states:
“Physical and chemical properties of the elements are the periodic function of their atomic numbers”.
The atomic mass of an element is determined by the mass of protons and neutrons in the nucleus of its atom. Since the nucleus is located inside an atom, it is not significantly linked with the properties of the element, particularly the chemical properties. These are related to the number of electrons and their distribution in the different energy shells. The elements with different electronic arrangements of atoms possess different chemical properties. As the number of electrons in an atom is given by the atomic number and not by the mass number, the atomic number should form the basis of the classification of elements in the periodic table, not the atomic mass as predicted by Mendeleev.
The repetition of similar properties of elements placed in a group and separated by a definite gap of atomic number is known as Periodicity.
The periodic properties can be defined as:
The properties of the elements are directly or indirectly related to the electronic configuration of their atoms and show gradation (increases or decreases) in moving down a group or a longer period.
The common physical properties of the elements include melting points, boiling points, density, enthalpy of fusion, and vaporization. However, we will focus mainly on the properties based on electronic configuration. These are:
The atomic radius is defined as the distance from the center of the nucleus to the outermost shell containing electrons. Depending on the type of bonding in the atoms, these are: (i) Covalent radii (ii) van der Waals radii (iii) Metallic radii
(i) Covalent radii: The covalent radius is defined as half the distance between the centers of the nuclei of two adjacent similar atoms joined to each other by a single covalent bond. For instance, the bond distance of Cl-Cl is 198 pm, making the covalent radius of Cl 99 pm.
(ii) van der Waals radii: The van der Waals radius is defined as half the internuclear distance between two similar adjacent atoms belonging to the two neighboring molecules of the same substance in the solid state.
(iii) Metallic radii: The metallic radius is defined as half the distance between the center of the nuclei of two adjacent atoms in the metallic crystal.
As we move from left to right in a period, the atomic radius decreases due to an increase in effective nuclear charge (Zeff). Along the group, as we move from top to bottom, atomic radius increases due to an increase in the principal quantum number, which leads to an increase in the number of shells, and an increase in shielding effect.
Ions are formed when a neutral atom loses or gains one or more electrons. An ion that loses electrons becomes a cation (positive ion), while an ion that gains electrons becomes an anion (negative ion). The ionic radius is defined as the effective distance from the center of the nucleus of the ion to the outermost part of the electron cloud.
The trends in ionic radii are the same as those of atomic radii. It decreases across a period from left to right and increases down a group. The size order of a cation, a neutral atom, and an anion of any natural atom is: cation < neutral atom < anion.
Ionization enthalpy is the amount of energy required to remove an electron from the outermost orbit of an isolated gaseous atom.
Generally, ionization energy increases from left to right across a period and decreases down a group. However, half-filled and fully filled orbitals are highly stable and thus have high ionization energy.
Electron gain enthalpy is the change in enthalpy that occurs when a gaseous atom gains an extra electron to form a monovalent anion in the gaseous state.
Electron gain enthalpy increases across periods and decreases down groups. Among all elements, chlorine has higher electron affinity than fluorine.
Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a covalent bond. Fluorine is the most electronegative element, while cesium is the least.
In periods, electronegativity increases from left to right. In groups, electronegativity decreases as we move down.
What is the need for the classification of elements?
Before the eighteenth century, only a few elements were known. It is quite easy to study and remember the properties of the element individually but at present, as many as 118 elements are known. It is impossible to remember the properties of each element and its compounds. Therefore many attempts have been made to classify elements into fewer groups; the purpose of classification has been to make the study of chemistry of elements and their compounds easier.
What is periodicity in the periodic table?
The periodicity in the properties of the elements placed in any group is due to the repetition of the same valency cell electronic configuration after a certain definite gap of atomic numbers(magic numbers) such as 2, 8, 8, 18, 18, 32.
What are 4 periodic properties?
The 4 periodic properties are atomic radii, ionization energy, electron affinity, electronegativity etc.
What is the basic classification of the modern periodic table?
In the modern periodic table, Physical and chemical properties of the elements are the periodic function of their atomic number.