Classification of Elements was necessary since many elements were being discovered in the 19th century and the study of these elements individually was proving difficult.
There were many attempts at classifying elements including ‘Dobereiner’s Triads’ and ‘Newland’s Octaves’.
German chemist Johann Wolfgang Dobereiner attempted to classify elements with similar properties into groups of three elements each. These groups were called ‘triads’. Dobereiner suggested that in these triads, the atomic mass of the element in the middle would be more or less equal to the mean of the atomic masses of the other two elements in the triad.
An example of such a triad would be one containing lithium, sodium, and potassium. The atomic mass of lithium 6.94 and that of potassium is 39.10. The element in the middle of this triad, sodium, has an atomic mass of 22.99 which is more or less equal to the mean of the atomic masses of lithium and potassium (which is 23.02).
The Limitations of Dobereiner’s Triads are :
English scientist John Newlands arranged the 56 known elements in increasing order of atomic mass in the year 1866. He observed a trend wherein every eighth element exhibited properties similar to the first. This similarity in the properties of every eighth element can be illustrated as follows.
Newland’s Law of Octaves states that when the elements are arranged in increasing order of atomic mass, the periodicity in properties of two elements which have an interval of seven elements in between them would be similar.
Limitations of Newland’s octaves are:
Russian chemist Dmitri Ivanovich Mendeleev put forth his periodic table in 1869. He observed that the properties of elements, both physical and chemical, were periodically related to the atomic mass of the elements.
The Periodic Law (also referred to as Mendeleev’s Law), states that the chemical properties of elements are a periodic function of their atomic weights.
The advantages of Mendeleev’s Periodic table are:
The limitations of Mendeleev’s Periodic table are:
These methods were the foundation on which the modern periodic table was built. However, the greatest contributor to the modern periodic table was Dmitri Mendeleev. Mendeleev is also known as the Father of the Modern Periodic Table. The modern periodic law is also called Mendeleev’s Law to honour him.
In the year 1913, English physicist Henry Moseley studied the wavelength of the characteristic x-rays By using different metals as anti cathode and showed that the square root of the frequency of the line is related to the atomic number. On the basis of the above observations Moseley gave the modern periodic law which states that :
“Physical and chemical properties of the elements are the periodic function of their atomic numbers”.
The atomic mass of an element is due to the mass of protons and neutrons present in the nucleus of its atom. Since the nucleus is located inside an atom, It is not very much linked with the properties of the element, particularly the chemical properties. These are related to the number of electrons and also the distributions of the electrons 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, therefore atomic number should form the basis of the classification of the elements in the periodic table and not atomic mass as predicted by Mendeleev.
Repetitions of the similar properties of the elements placed in a group and separated by certain definite gap of atomic number are known as Periodicity.
The periodic properties may 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 are melting points, boiling points, density, enthalpy of fusion and vaporization etc. But we shall focus our attention mainly on the properties which are based on electronic configuration these are:
The atomic radius may be defined as the distance from the centre of the nucleus to the outermost shell containing electrons. Depending upon the nature of bonding in the atoms these are (i) Covalent radii (ii) van der Waals radii (iii) Metallic radii
(i) Covalent radii: One-half of the distance between the centres of the nuclei of two adjacent similar atoms joined to each other by a single covalent bond is known as covalent radii. Eg Cl-Cl bond distance=198 pm covalent radius of Cl= 99 pm.
(ii)van der Waals radii: Half of the internuclear distance between two similar adjacent atoms belonging to the two neighbouring molecules of the same substance in the solid state is known as van der Waals radii.
(iii) Metallic radii: Half the distance between the centre of the nuclei of two adjacent atoms in the metallic crystal is known as metallic radii
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 increase in principal quantum number which causes an increase in the number of shells and increases in shielding effect .
The ions formed by the loss of one or more electrons from the neutral atom are known as cation (positive ion) when the electrons added to the neutral atom form an anion (negative ion). The effective distance from the centre of the nucleus of the ion upto which it exerts its influence on the electron cloud is known as the ionic radii.
The ionic radii change in the same trend as atomic radii. It decreases along the period from left to right and increases down the group from top to bottom. size of cation and anion of any natural atom as: cation< neutral atom < anion
The amount of energy required when an electron is removed from the outermost orbit of an isolated gaseous atom is known as Ionisation Enthalpy (IE).
Generally left to right in period IE increases whereas on moving down the group it decreases but half-filled orbital and fully filled orbitals are highly stable and thus have high IE.
The electron gain enthalpy is defined as the change in enthalpy which takes place when a gaseous atom gains an extra electron to form a monovalent anion in the gaseous state.
Electron gain enthalpy increases across the periods while it decreases down the group.
Chlorine has the highest 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 the periods left to right electronegativity increases. In the groups while moving down the groups electronegativity decreases .
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.
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.
The 4 periodic properties are atomic radii, ionization energy, electron affinity, electronegativity etc.
In the modern periodic table, Physical and chemical properties of the elements are the periodic function of their atomic number.
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