• In 1829, Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and the third elements.
  • A few triads proposed by him are listed.


The triads given by Dobereiner were helpful in grouping some elements with similar characteristics together, but he could not arrange all the elements known at that time into triads.


  • John Newlands proposed the law of octaves by stating that when elements are arranged in order of increasing atomic masses, every eighth element has properties similar to the first.
  • Newlands called it the law of octaves because a similar relationship exists in the musical notes also.
  • This can be illustrated as:


(i) This classification was successful only up to the element calcium. After that, every eighth element did not possess the same properties as the element lying above it in the same group. 
(ii) When noble gas elements were discovered at a later stage, their inclusion in the table disturbed the entire arrangement. 


  • Mendeleev’s Periodic Law: The physical and chemical properties of the elements are a periodic function of their atomic masses.
  • Mendeleev arranged the elements known at that time in order of increasing atomic masses and this arrangement was called the periodic table.
  • Elements with similar characteristics were present in vertical rows called groups. The horizontal rows were known as periods.


(i) In the periodic table, the elements are arranged in vertical rows called groups and horizontal rows known as periods. 
(ii) There are nine groups indicated by Roman Numerals as I, II, III, IV, V, VI, VII, VIII and zero. Group VIII consists of nine elements which are arranged in three triads. The zero group contains elements belonging to inert gases or noble gases and elements present have zero valency. 
(iii) There are seven periods (numbered from 1 to 7) or, horizontal rows in the Mendeleev’s periodic table. 


(i) This made the study of the elements quite systematic in the sense that if the properties of one element in a particular group are known, those of others can be predicted. 
(ii) This helped to a great extent in the discovery of these elements at a later stage. 
(iii) Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and properties. 


(i) Hydrogen has been placed in group IA along with alkali metals. But it also resembles halogens of group VII A in many properties.Thus, its position in Mendeleev’s periodic table is controversial. 
(ii) Although the elements in Mendeleev’s periodic table have been arranged in order of their atomic masses, in some cases the element with higher atomic mass precedes the element with lower atomic mass. 
(iii) We know that the isotopes of an element have different atomic masses but the same atomic number. 
Since, periodic table has been framed on the basis of increasing atomic masses of the elements, different positions must have been allotted to all the isotopes of a particular element. 
(iv) According to Mendeleev, the elements placed in the same group must resemble their properties. But there is no similarity among the elements in the two sub-groups of a particular group.                                                                        (v) In some cases, elements with similar properties have been placed in different groups.                                                                                                              (vi) Lanthanoids and actinoids were placed in two separate rows at the bottom of the periodic table without assigning a proper reason.                      (vii) No proper explanation has been offered for the fact that why the elements placed in group show resemblance in their properties.



  • Physical and chemical properties of the elements are the periodic function of their atomic numbers.
  • Present Form of the Periodic Table (Long form of Periodic Table)
    The long form of periodic table, also called Modem Periodic Table, is based on Modern periodic law. In this table, the elements have been arranged in order of increasing atomic numbers.
  • Nomenclature of Elements with Atomic No. more than 100


  • The long form of the periodic table also consists of the vertical rows called groups.
  • There are in all 18 groups in the periodic table.
  • Unlike the Mendeleev periodic table, each group is an independent group.


(i) All the elements present in a group have the same general electronic configuration of the atoms. 
(ii) The elements in a group are separated by definite gaps of atomic numbers (2, 8, 8,18, 18,32). 
(iii) The atomic sizes of the elements in the group increase down the group due to an increase in the number of shells. 
(iv) The physical properties of the elements such as m.p., b.p. density, solubility etc., follow a systematic pattern. 
(v) The elements in each group have generally similar chemical properties. 


  • Horizontal rows in a periodic table are known as periods.
  • There are seven periods in the long form of the periodic table.
  • Characteristics of periods:
    (i) In all the elements present in a period, the electrons are filled in the same valence shell.
    (ii) The atomic sizes generally decrease from left to right.


General electronic configuration: ns1-2 Characteristics of s-block elements:
(i) All the elements are soft metals. 
(ii) They have low melting and boiling points. 
(iii) They are highly reactive. 
(iv) Most of them impart colours to the flame. 
(v) They generally form ionic compounds. 
(vi) They are good conductors of heat and electricity. 


  • General electronic configuration: ns2np1-6
  • Characteristics of p-block elements:
    (i) The compounds of these elements are mostly covalent in nature.
    (ii) They show variable oxidation states.
    (iii) In moving from left to right in a period, the non-metallic character of the elements increases.
    (iv) The reactivity of elements in a group generally decreases downwards.
    (v) At the end of each period is a noble gas element with a closed valence shell ns2 np6 configuration.
    (vi) Metallic character increases as we go down the group.


  • General electronic configuration: (n -1) d1-10 ns0-2
  • The d-block elements are known as transition elements because they have incompletely filled d-orbitals in their ground state or in any of the oxidation states.

    (i) They are all metals with high melting and boiling points.                          (ii) The compounds of the elements are generally paramagnetic in nature.                                                                                                                      (iii) They mostly form coloured ions, exhibit variable valence (oxidation states).                                                                                                                    (iv) They are oftenly used as catalysts.


  • General electronic configuration: (n – 2) f1-14 (n -1) d0-1 ns2
  • They are known as inner transition elements because in the transition elements of d-block, the electrons are filled in (n – 1) d sub-shell while in the inner transition elements of f-block the filling of electrons takes place in (n – 2) f subshell, which happens to be one inner subshell.

    (i) The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids Ce (Z = 58) – Lu (Z = 71) and Actinoids Th (Z = 90) – Lr (Z = 103).                                                                                                                        (ii) These two series of elements are called Inner Transition Elements (f-Block Elements).                                                                                                    (iii) They are all metals. Within each series, the properties of the elements are quite similar.                                                                                (iv) Most of the elements pf the actinoid series are radio-active in nature.


(i) Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. 
(ii) Metals are solids at room temperature. 
(iii) Metal usually has high melting and boiling points. 
(iv) They are good conductors of heat and electricity. 
(v) They are malleable and ductile. 


(i) Non-metals are located at the top right hand side of the Periodic Table. 
(ii) Non-metals are usually solids or gases at low temperature with low melting and boiling points. 
(iii) They are poor conductors of heat and electricity. 
(iv) The non-metallic character increases as one goes from left to right across the Periodic Table. 
(v) Most non-metallic solids are brittle and are neither malleable nor ductile. 


  • The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) show the characteristic of both metals and non-metals.
  • These elements are also called semimetals.


  • These are the elements present in group 18.
  • Each period ends with a noble gas element.
  • All the members are of gaseous nature and because of the presence of all the occupied filled orbitals, they have very little tendency to take part in chemical combinations.
  • These are also called inert gases.


The elements of group 1 (alkali metals), group 2 (alkaline earth metals) and group 13 to 17 constitute the representative elements. They are elements of s-block and p-block.


The transition elements include all the d-block elements and they are present in the centre of the periodic table between s and p-block elements.


  • Lanthanide (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements.
  • They are also called f-block elements.
    The elements after uranium are also called transuranic elements.



  • Lanthanide (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements.
  • They are also called f-block elements.
  • The elements after uranium are also called transuranic elements.

Atomic Radii

  • It is defined as the distance from the centre of the nucleus to the outermost shell containing the electrons.
  • Depending upon whether an element is a non-metal or a metal, three different types of atomic radii are used.
  • These are:
    (i) Covalent radius
    (ii) Ionic Radius
    (iii) van der Waals radius
    (iv) Metallic radius.
(i) Covalent Radius

It is equal to half of the distance between the centres of the nuclei of two atoms held together by a purely covalent single bond.

(ii) Ionic Radius
It may be defined as the effictive distance from the nucleus of an ion upto which it has an influence in the ionic bond.
(iii) van der Waals radius
Atoms of Noble gases are held together by weak van der Waals forces of attraction. The van der Waals radius is half of the distance between the centre of nuclei of atoms of noble gases.
(iv) Metallic Radius
It is defined as half of the intemuclear distance between the two adjacent metal ions in the metallic lattice.

Variation of Atomic Radius in the Periodic Table

Variation in a Period: Along a period, the atomic radii of the elements generally decreases from left to right.
Variation in a group: The atomic radii of the elements in every group of the periodic table increases as we move downwards.

Ionic Radius

The ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals. In general, the ionic radii of elements exhibit the same trend as the atomic radii.


The removal of an electron from an atom results in the formation of a cation. The radius of cation is always smaller than that of the atom.


Gain of an electron leads to an anion. The radius of the anion is always larger than that of the atom.

Isoelectronic Species

Some atoms and ions which contain the same number of electrons, we call them isoelectronic species. For example, O2-, F–, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges.

Ionization Enthalpy

  • It is the energy required to remove an electron from an isolated gaseous atom in its ground state.
    M (g) + I.E ——->M+ (g) + e–
  • The unit of ionization enthalpy is kJ mol-1 and the unit of ionization potential is electron volt per atom.
  • Successive Ionization Enthalpies:
    If a gaseous atom is to lose more than one electron, they can be removed one after the other i.e., in succession and not simultaneously. This is known as successive ionization enthalpy (or potential).

Variation of Ionization Enthalpies in the Periodic Table

Variation of Ionization Enthalpy Along a Period

Along a period ionization enthalpies are expected to increase in moving across from left to the right, because the nuclear charge increases and the atomic size decreases.

Variation of Ionization Enthalpy in a Group

  • The ionization enthalpies of the elements decrease on moving from top to the bottom in any group.
  • The decrease in ionization enthalpies down any group is because of the following factors.
    (i) There is an increase in the number of the main energy shells (n) in moving from one element to the other.
    (ii) There is also an increase in the magnitude of the screening effect due to the gradual increase in the number of inner electrons.

Electron Gain Enthalpy

Electron Gain Enthalpy is the energy released when an electron is added to an isolated gaseous atom so as to convert it into a negative ion. The process is represented as:
  • For the majority of the elements the electron gain enthalpy is negative.
  • For example, the electron gain enthalpy for halogens is highly negative because they can acquire the nearest noble gas configuration by accepting an extra electron.
  • In contrast, noble gases have large positive electron gain enthalpies because the extra electron has to be placed in the next higher principal quantum energy level thereby producing a highly unstable electronic configuration.

Successive Electron Gain Enthalpies:

  • We have studied that electrons from gaseous atoms are lost in succession (i.e., one after the other).
  • Similarly, these are also accepted one after the other, i.e., in succession.
    After the addition of one electron, the atom becomes negatively charged and the second electron is to be added to a negatively charged ion.
  • But the addition of a second electron is opposed by electrostatic repulsion and hence the energy has to be supplied for the addition of the second electron.
  • Thus the second electron gain enthalpy of an element is positive.
    For example, when an electron is added to an oxygen atom to form A– ion, energy is released.
  • But when another electron is added to 0- ion to form O2- ion, energy is absorbed to overcome the strong electrostatic repulsion between the negatively charged 0– ion and the second electron being added. Thus, the first electron gain enthalpy:

Factors on which Electron Gain Enthalpy Depends:

(i) Atomic size: As the size of an atom increases, the distance between its nucleus and the incoming electron also increases and electron gain enthalpy becomes less negative, 
(ii) Nuclear charge: With the increase in nuclear charge, force of attraction between the nucleus and the incoming electron increases and thus electron gain enthalpy becomes more negative. 
(iii) Symmetry of the Electronic Configuration: The atoms with symmetrical configuration (having fully filled or half filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unstable.  In that case the energy will be needed and electron gain enthalpy (Δ eg H) will be positive. For example, noble gas elements have positive electron gain enthalpies. 

Variation of Electron Gain Enthalpy Across a Period:

  • Electron gain enthalpy becomes more negative with an increase in the atomic number across a period.
  • Variation of Electron Gain Enthalpy in a Group:
    Electron gain enthalpy becomes less negative as we go down a group.


  • A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity.
  • Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity.
  • However, a number of numerical scales of electronegativity of elements viz, Pauling scale, Milliken- Jaffe scale, Allred Kochow scale have been developed.
  • The electronegativity of any given element is not constant; it varies depending on the element to which it is bound.
  • Across a Period:
    Electronegativity generally increases across a period from left to right.
  • In a Group:
    It decreases down a group.

Periodic Trends in Chemical Properties along a Period

(i) Metallic character: Decrease across a period maximum on the extreme left (alkali metals). 
(ii) Non-metallic character: Increasess along a period. (From left to right). 
(iii) Basic nature of oxides: Decreases from left to right in a period. 
(iv) Acidic nature of oxides: Increases from left to right in a period. 

Variation from Top to Bottom on Moving Down a Group

(i) Metallic character. Generally increases because of increase in atomic size and hence decrease in the ionizatiort energy of the elements in a group from top to bottom. 
(ii) Non-metallic character. Generally decreases down a group. As electronegativity of elements decreases from top to bottom in a group, 
(iii) Basic nature of oxides. Since the metallic character or electropositivity of elements increases when going from top to bottom in a group, the basic nature of oxidation naturally increases. 
(iv) Acidic character of oxides. Generally decreases as the non-metallic character of elements decreases in going from top to bottom in a group. 
(v) Reactivity of metals. Generally increases down a group. Since the tendency to lose electrons increases. 
(vi) Reactivity of non-metals. Generally decreases down the group, Higher the electro-negativity of non-metals, greater is their reactivity. Since electronegativity of non-metals in a group decreases from top to bottom, their reactivity also decreases. 

Anomalous Properties of Second Period Elements

  • The first element of each of the group 1 (lithium) and 2 (beryllium) and group 13-17 (boron to fluorine) differs in many respects from the other members of their respective groups.
  • For example, lithium unlike other alkali metals, and beryllium unlike other alkaline earth metals form compounds which have significant covalent character; the other members of these groups, pre-dominatly form ionic compounds.
  • It has been observed that some elements of the second period show similarities with the elements of the third period placed diagonally to each other, though belonging to different groups.
  • For example,
This similarity in properties of elements placed diagonally to each other is called a diagonal relationship.

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