What is the Periodic Table?

• The periodic table is an arrangement of all the elements known to man by their increasing atomic number and recurring chemical properties. 

• They are assorted in a tabular arrangement wherein a row is a period and a column is a group. Elements are arranged from left to right and top to bottom in the order of their increasing atomic numbers. 

• Thus,

  • Elements in the same group will have the same valence electron configuration and hence, similar chemical properties.

  • Whereas, elements in the same period will have an increasing order of valence electrons. 

    • Therefore, as the energy level of the atom increases, the number of energy sub-levels per energy level increases.

• The first 94 elements of the periodic table are naturally occurring, while the rest from 95 to 118 have only been synthesized in laboratories or nuclear reactors. 

  • The modern periodic table, the one we use now, is a new and improved version of certain models put forth by scientists in the 19th and 20th century. 

• Dimitri Mendeleev put forward his periodic table based on the findings of some scientists before him like John Newlands and Antoine-Laurent de Lavoisier. 

  • However, Mendeleev is given sole credit for his development of the periodic table.

Acceptance of Mendeleev Periodic Table:

• Dimitri Mendeleev, widely referred as the father of the periodic table put forth the first iteration of the periodic table similar to the one we use now. 

  • Mendeleev’s periodic law is different from the modern periodic law in one main aspect.

• Mendeleev modeled his periodic table on the basis of increasing atomic mass, whereas, the modern periodic law is based on the increasing order of atomic numbers.

• Even though Mendeleev’s periodic table was based on atomic weight, he was able to predict the discovery and properties of certain elements. 

• During his time only around half of the elements known to us now were known, and most of the information known about the elements were inaccurate. 

  • Mendeleev’s Periodic Table was published in the German Journal of chemistry in 1869.

Genesis of Periodic Classification:

Dobereiner’s Trend:

• 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 is 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:

  • All the elements known at that time couldn’t be classified into triads.

  • Only four triads were mentioned – (Li,Na,K ),  (Ca,Sr,Ba) ,  (Cl,Br,I) , (S,Se,Te).

Newland’s Octaves:

• 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:

  • It was only up to calcium that the classification of elements was done via Newland’s Octaves.

  • The discovery of noble gases added to the limitations of this method since they couldn’t be included in this arrangement without disturbing it completely.

Mendeleev’s Periodic Table:

• 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 inclusion of these newly discovered elements did not disturb the periodic table. Examples include germanium, gallium, and scandium.

  • It was used to correct the wrong atomic weights in use at that time.

  • A variance from the atomic weight order was provided by Mendeleev’s table.

• The limitations of Mendeleev’s Periodic table are:

  • Hydrogen’s position was in the group of alkali metals but hydrogen also exhibited halogen like qualities.

  • Isotopes were positioned differently since this type of classification of elements was done by considering the atomic weight of the element. 

    • Therefore – protium, deuterium, and tritium would occupy varying positions in Mendeleev’s table.

  • An anomalous positioning of a few elements showed that the atomic masses did not increase regularly from one element to the next. 

    • An example of this would be the placement of cobalt (atomic mass of 58.9) before nickel (atomic mass of 58.7).

• 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.

Modern Periodic Table:

• 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.

Classification of elements in modern periodic table

• The modern periodic table consists of 18 vertical columns, called the groups(1-18) and 7 Horizontal rows, called periods.

  • The first period contains two elements, Hydrogen and Helium.

  • The second period contains eight elements, from Lithium to Neon.

  • The third period contains eight elements, from Sodium to Argon.

  • The fourth period contains eighteen elements, from Potassium to Krypton.

  • The fifth period contains eighteen elements, from Rubidium to Xenon.

  • The sixth period contains thirty-two elements.

  • The seventh period is incomplete.

• On the basis of electronic configuration, elements are classified into four Blocks known as s, p, d and f- blocks.

  • 1st and 2nd group elements are called s-block elements. The general electronic configuration is ns1-2.

  • 13th to 18th group elements are called p-block elements. The general electronic configuration is ns2 np1-6.

  • 3rd to 12th group elements are called d-block elements. The general electronic configuration is (n-1)d1-10 ns1-2.

  • Lanthanides and actinides elements are called f-block elements. The general electronic configuration is (n-2)f1-14 (n-1)d0-1 ns2.

• Periodic properties and their trends

  • 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:

    • Atomic and ionic radii 

    • Ionization enthalpy

    • Electrons gain enthalpy

    • Electronegativity

Atomic and Ionic Radii:

Atomic Radii:

• 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:

  • 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.

  • 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.

  • 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.

Ionic Radii:

• 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 

• Ionization enthalpy:

  • 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.

Electrons gain enthalpy:

• 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:

• 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.