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Periodic Table

Periodic table  is an arrangement of elements with similar properties placed together. The periodic table evolved largely as a result of experimental observations.

Earlier Attempts to Classify Elements:

(1) Dobereiner’s law of triads (1829) was the classification of elements into groups of three elements each with similar properties such that the atomic weight of the middle element was the arithmetic mean of the other two e.g. Ca, Sr, Ba, Cl, Br, I etc

(2) Telluric screw or Helix was proposed by Chancourtois in 1862.

       (3) Newlands law of octaves (1863) was an arrangement of elements in order of increasing atomic weights in which it was observed that every eighth element had properties similar to those of the first just like the eighth note of an octave of music.

(4) Mendeleef’s Periodic Table 

Mendeleef arranged the then discovered 63 elements in the periodic table into 7 horizontal rows known as periods and 18 vertical columns known as groups numbered 1 to 18.

Mendeleefs Periodic law

(i)        The physical and chemical properties of elements are periodic function of their atomic weights

(ii)       If the elements are arranged in the order their increasing atomic weights, after a regular interval, elements with similar properties are repeated (inert gases were not discovered then)

Mendeleef Suggested

(i)        Exclusion of certain elements and assigned them a separate independent position.

(ii)       Leaving gaps for the then undiscovered elements

(iii)      When the properties of elements did not correspond to what is expected of the group they were named by prefixing Eka to the preceding elements e.g. Eka boron (Silicon); Eka silicon (Germanium), Eka aluminum (Gallium); Eka Maganese (Technitium)

Uses of Mendeleefs Periodic Table

(i)        Atomic weights of elements were determined with the help of periodic table.

Atomic weight = Valency x  Equivalent wt. = Group number x  Equivalent weight.

(ii)       Atomic weight of elements were corrected. Atomic weight of Be was calculated to be 3 x4.5 = 13.5  by considering its valency Mendeleef calculate it 2 x 4.5 = 9

(iii)      Discovery of new elements–In Mendeleef’s periodic table two consecutive members differ by two or three units in the atomic weight Where this gap was more, the gaps were left in the periodic table.

Defects of Mendeleef’s Periodic Table  

(i)        Position of hydrogen is uncertain. It has been placed in A and VIIA groups because of its resemblance with both groups.

(ii)       No separate positions given to isotopes

(iii)      It is not clear to which group lanthanides and actinides belong to.

(iv)      Although there is no resemblance except valency of subgroups A and B, they have been put in the same group.

(v)       Order of increasing atomic weight is not strictly followed in the arrangement of elements in the periodic table. For e.g. – Co (At. wt. 58.9) is placed before I(127) and Ar (39.9) before K(39)

Long form of the Periodic Table or Moseley’s Periodic Table

(i)        Mosely (1909) studies the frequency of X-rays produced by the bombardment of a strong beam of electrons on a metal target. He found that the square root of the frequency of Xray (√v) is directly  proportional to the number of nuclear charge (z) of metal.                      √v = a (Z-b)where a and b are constants. Nuclear charge of metal is equal to the atomic number. So Moseley related the properties of elements with their atomic number and gave the new periodic law.

(ii)       Moseley’s Periodic Law or Modern Law: Physical and chemical properties of elements are the periodic function of their atomic numbers. If the elements are arranged in order of their increasing atomic number, after a regular interval, elements with similar properties are repeated.

Characteristics of Periodic  

(i)        First period is called shortest period and contains only two elements. Second and third periods are called short periods containing eight elements each. Sixth period is longest period with thirty-two elements seventh period is an incomplete period.

(ii)       Lanthanide and actinide series containing 14 elements each are placed separately under the main periodic table. These are related to sixth and seventh periods of III group respectively

(iii)      Elements of third period from sodium (Na) to Chlorine (Cl) are called representative or typical elements

(iv)      Valency of an element in a period increases from 1 to 7 with respect to oxygen.

             Na2O     MgO    Al2O3    SiO2      P2O5     SO3    Cl2O7

              1            2            3            4            5           6          7

(v)       Elements of second period are called bridge elements.

Characteristic of Groups  

(i)        Moseley’s periodic table contains 18 groups. These are represented by roman numerals I, II, III, IV, V, VI, VII VIII and zero. Groups I to VII are divided into two subgroups A and B Group VIII consists of three sets, each one containing three elements.

(ii)       Inert gases are present in zero group.

(iii)      The Valency of an elements in a group is equal to the group number.

(iv)      There is no resemblance in the elements of subgroups A and B same group, except valency

(v)       The elements of the group which resemble with typical elements are called normal elements. For example IA, IIA, IIIA, IVA, VA, VIA, VIIA  group elements are normal elements.

(vi)      Those elements of the groups which do not resemble with typical elements are called transition elements. Fr example- IB, IIIB, IVB, VB, VIB, VIIB, and VIII group elements are transition.

Long From of the periodic Table and Electronic Configuration of Elements

(i)        Each period starts with an alkali metal whose outermost electronic configuration is ns1.

(ii)       Each period ends with a noble gas of outermost electronic configuration ns2np6 except He. The electronic configuration of He is 1s2

(iii)      The number of elements in a period is equal to the number of necessary electrons to acquire ns2np6 configuration in the outermost shell of first element (alkali metal) of the period. First period contains two elements.

(iv)      The number of elements in each period may be determined by the number of electrons in a stable configuration as under

Periods Stable electronic configuration Number of electrons
First 1s2 2 He
Second 2s22p6 8 Ne(10)
Third 3s23p6 8 Ar(18)
Fourth 4s23d104p6 18 Kr(36)
Fifth 5s24d105p6 18 Xe(54)
Sixth 6s24f145d106p6 32 Rn (86)


Classification of Elements on the Basis of their Electronic Configuration

On the basis of electronic configuration, the elements may be divided into four groups.

(i)        s-Block elements

            (a) These are present in the left part of the periodic table.

(b) These are IA and IIA i.e. 1 and 2 group elements

(c) These are metals

(d) In these elements last electron fills in the s-orbital

(d) Electronic configuration of valence shell is ns1-2 (n=1 to 7)

(ii)       p-Block elements

  • These are present in right part of the periodic table.
  • These constitute the group IIIA to VIIA and zero groups i.e. groups 13 to 18 of the periodic table
  • Most of these elements are metalloids and non metals but some of them are metals also.
  • The last electron fills in p-orbital of valence shell.
  • The electronic configuration of valence shell is ns2 np1-6 (n = 2 to 7)
  • ns2 np6 is stable noble gas configuration. The electronic configuration of He 1s2.

(iii)      d-Block Elements

  • These are present in the middle part of the periodic table (between s & p block element)
  • These constitute the group IIIB to VIIIB and IIB i.e. 3 to 12 groups of the periodic table.
  • All are metals.
  • The last electrons fills in (n–1) d orbital
  • The outermost electronic configuration is (n-1)d1-10ns1-2(n=4 to 7)
  • There are three series of d-block elements as under

3d series – Sc(21) to Zn (30)

4d series – Y (39) to Cd (48)

5d series – La (57), Hf (72) to Hg (80)

(iv)      f-Block Elements

  • These are placed separately below the main periodic table
  • These are mainly related to IIIB i.e. group 3 of the periodic table
  • There are two series of f-block element as under 4f series –Lanthanides – 14 Elements Ce (58) to Lu(71)

5f series – Actinides – 14 Elements Th (90) to Lw (103)

  • The last electron fills in (n-2) f orbital
  • Their outermost electronic configuration is (n-2)f1-14(n-1)p6(n-1)d0-1ns2(n = 6 and 7)

Bohr’s Classification of Elements

Elements may be classified into four groups on the basis of number of incomplete shells present in them-

(i)        Inert gases

(a) s-and p-orbitals on the outer most shell of these elements are completely filled. The outermost electronic configuration is ns2np6.

(b) Helium is also inert gas its electronic configuration is 1s2

(ii)       Representative or Normal Elements

  • Outermost shell of these elements is incomplete. The number of electrons in the outermost shell is less than eight.
  • Inner p-block elements except inert gases are called normal or representative elements.
  • S-and p-block elements except inert gases are called normal or representative elements.

(iii)      Transition Elements

  • Last two shells of these elements namely outermost and penultimate shells are incomplete
  • The last shell contains one or two electrons and penultimate shell may contain more than eight up to eighteen electrons
  • Their outermost electronic configuration is similar to d-block elements i.e. (n–1)1-10 ns1-2.
  • According to latest definition of transition elements those elements which have partly filled d-orbital in natural state or in any stable oxidation state are called transition elements but grp. IB & IIB are not transition elements because these elements have d10 configuration in neutral as well as in stable +2 oxidation state.

(iv)      Inner Transition Elements

  • In these elements last three shell i.e. last, penultimate and per penultimate shells are incomplete
  • These are related to IIIB i.e. group 3.
  • Theses last shell contains two electrons. Penultimate shell may contain eight or nine electrons and per penultimate shell contains more than 18 up to 32 electrons.
  • Their outmost electronic configuration is similar to f – block elements i.e
  • Elements of the seventh period after atomic number 92 (i.e actinides) are synthetic elements and are called transition elements

Periodicity in Atomic properties

(i)        In a period from left to right there, is a regular change in electronic configuration of elements. In a group from top 10 bottom the outermost shell electronic configuration is similar.

(ii)       The chemical properties of the elements depends upon their electronic configuration. So there is a regular change in chemical properties in a period and elements have similar chemical properties in a group.

(iii)      In a period as well as in group there is a regular gradation (gradual increase or decrease in particular property) in physical and chemical properties of elements with the change in atomic number. This regularity in properties is called periodicity.

Periodicity in some atomic properties is as under –

Atomic Radius     

(i)        The radius of an atom may be taken as the distance between atomic nucleus and the outermost shell of electrons of the atom.

(ii)       According to the Heisenberg’s uncertainty principle the position of a moving electron can not be accurately determined. So the distance between the nucleus and the outermost electron is uncertain.

(iii)      Atomic radius can be determined indirectly from inter nuclear distance between the two atoms in a gaseous diatomic molecule. This internu clear distance between the two atoms is called bond length.

(iv)      The inter nuclear distance between the two atoms can be measured by X – ray diffraction or spectroscopic studies.

(v)       Covalent radius – One half of the distance between the nuclei (inter nuclear distance) of two covalently bonded like atoms in a homo nuclear diatomic molecule is called the covalent radius of that. The covalent radius (rA) of atom A in a molecule A2 may be given as            – rA =dA-A/2 i.e. the distance between nuclei of two single covalently bonded atoms in a homo diatomic molecules is equal to the sum of covalent radii of both the atoms. dA-A=rA+rA  In a hetero diatomic molecules AB where the electro negativity of atoms A and B are different, the experimental values of inter nuclear distance DA-B is less than the theoretical value (rA+rB)  Aaccording to Shoemaker and Stevenson (1941) \(\displaystyle {{D}_{{A-B}}}={{r}_{A}}+{{r}_{B}}-0.09{{\Delta }_{x}}\) Where \(\displaystyle {{\Delta }_{x}}\) is the difference of electro negativities of the atoms of A and B.

According to Puling – If the electro negativities of the two atoms A and B are xA & xB  respectively then

\(\displaystyle {{D}_{{A-B}}}={{r}_{A}}+{{r}_{B}}-\left( {{{C}_{1}}{{x}_{A}}-{{C}_{2}}{{x}_{B}}} \right)\) C1 and C2 are the Stevenson’s  coefficients for atoms A and B respectively

(vi)      Metallic radius

Metal atoms are assumed to be closely packed spheres in the metallic crystal. These metal atom spheres are consided to touch one another in the crystal. One half of the inter nuclear distance between the two closed metal in the metallic crystal is called metallic radius.

Metallic radius > Covalent radius

For example – Metallic radius and covalent radius of potassium are 2.3 A and 2.03 A respectively

(vii)     Van der Waal’s Radius or Collision radius

The molecules of non metal atoms are generally gasses. On cooling, the gaseous state change to solid state. In the solid state, the non metallic elements usually exist as aggregations of molecules which are held together by Van der Waal’s force One half of the distance between the nuclei of two adjacent atoms belonging to two neighboring molecules of an elements in  the solid state is called Van der Wall’s radius

It may also be defined as half of the inter nuclear distance of two non bonded neighboring atoms of two adjacent molecules.

Vander Waal’s radius > Metallic radius > Covalent radius

The Vander Waal’s radius and covalent radius of Chlorine atom are 1.80Å and 0.99Å respectively.

Ionic Radius

A neutral atom changes to ion by the gain of one or more electrons. The number of charge on an ion is equal to the number of electrons lost or gained. The ionic radii of the ions present in an ionic crystal may be calculated from the inter nuclear distance between the two ions

(i)        Radius of a Cation – Radius of a cation is invariably smaller than that of the corresponding neutral atom

Na                   Na+

Number of e =                                               11                    10

Number of p =                                                11                    11


  • The affective nuclear charge increases. For example in Na atom 11 electrons are attracted by 11 protons and in Na+ 10 electrons are attracted by 11 protons. Thus in the formation of cation of electrons decreases and nuclear charge remains the same.
  • Generally the formation of cation results in the removal of the whose outer shell
  • Inter electronic repulsion decreases. The Inter electronic repulsion in Na is among 11e and in Na+ among 10e

(ii)       Radius of an anion –:    Radius of an anion is invariably bigger than that of the corresponding atom

Cl                    Cl–

Number of e  =          17                    18

Number of p    =         17                    17           


(a) The effective nuclear charge decrease in the formation of anion. Thus the electrostatic force of attraction between the nucleus and the outer electrons decreases and the size of the anion increases.

(b) Electronic series – A series of atoms, ions and molecules in which each specie contains same number of electrons but different nuclear charge is called is electronic series

N3–                  O2–                  F                    Ne                   Na+                        Mg2+

Number of e  10                    10                    10                    10                    10                        10

Number of p    7                      8                      9                      10                    11                        12

(a) Number of electrons is same.

(b) Number of protons is increasing.

(c)       So the effective nuclear charge is increasing and atomic size is decreasing. In an isoelectronic series atomic size decreases with the increase of charge. Some of the examples of isoelectronic series are as under

S2–, Cl, K+, Ca2+, Sc3+,           N2, CO, CN,              NH3, H3O+

Periodicity in Atomic Radius and Ionic Radius 

  1. For normal elements

(a) In a period from left to right, effective nuclear charge increases because the next electron fills in the same shell. So the atomic size decreases. For example the covalent radii of second period elements in are as follows–

Li                    Be                    B                     C                     N                     O                        F

1.23                 0.89                 0.80                 0.77                 0.74                 0.74                        0.72

(b)In a group moving from top to bottom the number of shells increases. So the atomic size increases. Although the effective nuclear charge increases but its effect is negligible in comparison to the effect of increasing number of shells. For example the covalent radii of IA group elements inare as follows-

Li                    Na                   K                     Rb                   Cs

1.23                 1.57                 2.03                 2.16                 2.35

  1. The atomic radius of inert gas (zero group) is shown largest in a period because of its Vander Waal’s radius which is generally larger than covalent radius. The Vander Waal’s radius of inert gases also increases in moving from top to bottom in a group.
  2. For transition elements – There are three series of transition elements-

3d – Sc (21) to Zn (30)

4d – Y (39) to Cd (48)

5d – La (57), Hf (72) to Hg (80)

(a) From left to right in a period

(i)     The atomic size decreases due to the increase in effective nuclear charge.

(ii)    In transition elements, electron are filled in the (n-1)d orbital. These (n-1)d electron screen the ns electrons from the nucleus. So the force of attraction between the ns electrons and nucleus decreases.

This effect of (n-1)d electrons over ns electrons is called shielding effect or  screening effect. The atomic size increases due to shielding effect and balance the increases in size due to increase in nuclear charge to about 80%.

(iii)   Thus moving from left to right in a period, there is a very small decrease in size and it may be considered that size is almost the same.

(iv)   In the first transition series the atomic size slightly decreases from Sc to Mn because effect of effective nuclear charge is stronger than the shielding effect. The atomic size from the Fe to Ni almost remains the same because both the effects balance each other. The atomic size from Cu to Zn slightly increases because of shielding effect. The atomic radii of the elements of 3d transition series are as under.

Sc           Ti        V         Cr        Mn      Fe        Co       Ni        Cu       Zn

1.44        1.32     1.22     1.18     1.17     1.17     1.16     1.15     1.17     1.25

(v)    Inner transition elements – As we move along the lanthanide series, there is a decrease in atomic as well as ionic radius. The decreases in size is regular in ions. But not regular in atoms. This is called lanthanide contraction. The atomic radii in  are as under –

La           Ce        Pr        Nd       Pm       Sm       Eu        Gd

1.88        1.82     1.83     1.82     –           1.80     2.04     1.80

Tb           Dy       Ho       Er        Yb       Lu

1.78        1.77     1.76     1.75     1.94     1.73

There are two peaks one at Eu (63) and other Yb (70). This is due to the difference in  contribution in  metallic bonding Except Eu and Yb other lanthanides contribute three electrons in metallic bond formation. These two atoms contribute electrons in the bond formation leaving behind half filled and completely filled 4f- orbital respectively.

Cause of Lanthanide contraction – In lanthanide the additional electron enter in (n-2)f orbital. The mutual shielding effect of (n-2)f electrons is very little because the shape of f-sub shell is very much diffused. Thus the nuclear charge increases in the mutual shielding effect of (n-2) f electrons. The outer electrons are attracted more by the nucleus. Consequently the atomic and ionic radii decreases from La(57) to Lu(71)

This type of contraction also occurs in actinides. The jump in contraction between the consecutive elements in the actinides is greater than lanthanides. This is due to the lesser shielding of 5f-electrons which are therefore pulled more strongly by the nucleus.

In a group

(i)        The atomic radius of elements increases moving from transition series (3d) to second transition series (4d). This is due to the increases in number of shells with the increases in atomic number.

(ii)       The atomic radii of second (4d) and third (5d) transition series in a group is almost same except Y(39) and La(57). In third transition series, there are fourteen lanthanides in between La (57) of III B and Hf (72) of IV B groups, so the atomic radius of Hf (72) decreases much due to lanthanide contraction in lanthanides. The difference in the nuclear charge in the elements of a group in first and second transition series is + 18 units while this difference in second and third transition series + 32 units except Y (39) La(57). Due to the increase of + 32 units in the nuclear charge there is a sizable decrease in the atomic radius which balances the increase in number of shells.

So in a group moving from second to third transition series, the atomic radii of the elements almost remain the except IIIB. The same difference is about 0.02 Å.

Ionization Potential or ionization Energy

(i)     The amount of energy required to remove the most loosely bound electron of the outermost shell (i.e. the outermost electron) from one mole of an isolated gaseous atom of an element in its ground state to produce a cation is known as ionization energy of that elements.

(ii)    Because ionization energy is generally expressed in electron volts, so it is also  known as ionization potential.

(iii)   Energy required for the removal of first, second and third electron from the gaseous atom is called first, second and third ionization energy respectively

\(\displaystyle {{A}_{{\left( g \right)}}}+{{I}_{1}}\to {{A}_{{\left( g \right)}}}^{+}+{{e}^{-}},A_{{\left( g \right)}}^{+}+{{I}_{2}}\to A_{{\left( g \right)}}^{{2+}}+{{e}^{-}},A_{{\left( g \right)}}^{{2+}}{{I}_{3}}\to A_{{\left( g \right)}}^{{3+}}+{{e}^{-}}\)

(iv)   The order of first, second and third ionization energies may be given as l1 < l2 < l3

This is because second and third electron is removed from monopositive and dispositive cations respectively. Effective nuclear charge increases with the increase of positive charge. So the attraction between the nucleus and the outermost electron increases and more energy is required for the removal of electron.

Factors Affecting Ionization Potential

(i)     Atomic radius: The value of ionization potential of an element- decreases as its atomic radius increases. This is because the electrostatic force of attraction between the nucleus and the outermost electron decreases as the distance between them increases. So the energy required for the removal of electron will comparatively be less

\(\displaystyle Ionisation\,\,pdential\,\alpha \frac{1}{{Atomic\,radius}}\)

(ii)    Effective nuclear charge: The greater the effective charge on the nucleus of an atom, the more difficult it be to remove an electron from the atom because electrostatic force of attraction between the nucleus and the outermost electron increases, So the greater energy will be required to remove the electron.

         Ionization potential α Effective nuclear charge(Zeff)

(iii)   Penetration effect of orbitals: The order of energy required to remove electron from s,p,d-and f-orbitals of a shell is s> p> d>f because the distance of the electron from the nucleus increases. For example – The value of ionization potential of Be (z = 4, 1s22s2) and Mg(Z = 12, 1s2 2s22p63s2) are more than the I.P.s of B (Z=5, 1s22s22p1x) and Al(Z = 13, 1s22s22p63s23p1x) because the penetration power of 2s and 3s electrons is more than 2p and 3p orbitals respectively. More energy will be required to separate the electrons from 2s and 3s orbital.

(iv)   Shielding or screening effect: The shielding or screening effect increases if the number of electrons increases. This results in decrease of force of attraction between the electron. Thus the value of I.P decreases.

\(\displaystyle Ionisation\,\,potentrial\,\,\alpha \frac{1}{{Shielding\,\,effect}}\)

(v)    Stability of half –filled and fully- filled orbitals: According to Hund’s rule the stability of half filled and completely filled degenerate orbitals are comparatively high. So comparatively more energy is required to separate the electron from such atoms.

For example 

(a)       Removal of electron is comparatively difficult from the half filled configuration of N

\(\displaystyle \left( {Z=7,l{{s}^{2}}2{{s}^{2}}p{{x}^{1}}p{{y}^{1}}p{{z}^{1}}} \right)\)

(b)  The ionization potential of inert gases is very high due to stable s2p6 electronic  configurations.

Periodicity in Ionisation Potential 

(i)     For normal elements: On moving from left to right in a period, value of ionisation potential of elements increases because effective nuclear charge also increases.

Exceptions –

(a)   In a period, the ionisation energy of IIA group elements is more than the elements of IIIA because penetration power of s-electrons. The value of ionisation energy of Be(1s22s2) is more than B (1s22s2p1x)because the penetration power of 2s-electrons of Be is more than the 2px- electrons of B.

(b)  In a period, the ionisation energy of VA elements is more than the elements of VIA because the half filled p3 configuration of VA elements is comparatively of higher stability. VIA group elements (p4) have the tendency to acquire comparatively more stable (p3) configuration by the loss of one electron. Ionisation energy N(1s22s22p1xpy1pz1)>O 1s2 2s2 2p2xp1x p1x p1z Thus P > S, As> Se.

But the value of I.P of Sb (VA) & Te (VIA) and Bi (VA) are according to general rule i.e.

Sb (VA) < Te (VIA)           Bi(VA < Po(VIA)

On moving from top to bottom in a group the value of I.P decreases because the atomic size increases


(a)   In group IIIA the ionisation potential of Al (13) is equal to the ionisation potential of Ga (31). Before Ga(31) the electrons are filled in 3d-orbitals of ten transition elements. These 3d orbitals electrons do not completely shield the 4p- electron. So the increases of +18 units in nuclear charge result in the greater increase of effective nuclear charge. Due to increases in nuclear charge the I.P. increases which counter balance the decrease in I.P. due to the increase in number of shells.

(b)  The values of I.P. of Tl (81) and Pb(82) of sixth period is more than that I.P. values of In (49) and Sn (50) of same group in period fifth. This is because of the electrons are filled in 4f-orbitals before Tl (81) and Pb (82) which do not completely shied the outer electrons. Thus increase in +32 units in nuclear charge results in the increase of ionisation potential value.

(ii)    For transition elements: On moving from left to right in a transition series –

(a)   As the atomic number increases the effective unclear charge also increases hence the I.P. increases.

(b)  The shielding effect of (n–1) d electrons over ns electrons increases with the addition of electron in (n–1) orbitals. Hence the I.P. decreases.

(c)   The increase in values of I.P due to the increase of effective nuclear charge almost balance the decreases of I.P due to increase in shielding effect. There is a very small increase in the value of I.P or it may be said that I.P almost remains the same.

(d)  In first transition series from Sc to Cr the value of I.P increases because effect of increase ineffective nuclear charge is more than the shielding effect I.P values of Fe, Co, Ni and Cu are almost same. Due to d10s2 configuration of Zn, the first I.P. increases.

On moving from top to bottom in a group in transition series         

(a)    In a group on moving from first to second transition series, the values of I.P decreases because atomic size increases.

(b)       When moving from second to third transition series the value of I.P some what increases except IIIB group \(\displaystyle \left[ {Y\left( {39} \right)\to La\left( {57} \right)} \right].\)

This is because of 14 electrons are filed in 4f-orbitals of lanthanides which do not shield than 5d electrons effectively. Thus the increase in +32 units in nuclear charge results in the increase of I.P., on moving from left to right this effect decreases and becomes negligible in the later part.

Applications of lonization potential

(i)     Metallic or electropositive character of elements increases as the value of ionisation potential decreases. So in a group moving from top to bottom metallic or electropositive character increases because I.P value decreases. In a period moving from left to right the value of I.P increases so metallic or electropositive character decreases. non metallic character increases.

(ii)    The relative reactivity of the metals increases with the decrease in I.P values. The I.P. value of IA and IIA metals are comparatively low. So they are comparatively more reactive. The I.P. value of inert gases are very high. So they are almost inactive.

In a group moving from top to bottom the reactivity of metal atoms increases because their I.P. value decreases.

(iii)   The reducing power of elements increases as the value of I.P. decreases tendency to loss the electron increases. The reducing power increases going down group because the I.P. value decreases. Li is exception in IA group. The reducing power of Li is highest in its own group. The order of reducing power of IA elements is as under. Li > Cs > Rb > K > Na

(iv)   Determination of oxidation state or Valence electrons of an elements –

(a)   If the difference of two consecutive I.P.’s of an elements is 16 eV or more, the lower oxidation state is stable. For e.g. the difference of first and second I.P. of Na is 42.4 eV which is more than 16eV. So Na+ will be stable. It can also be explained from its electronic configuration

\(\displaystyle \begin{array}{l}Na\left( {11} \right)=1{{s}^{2}}2{{s}^{2}}2{{p}^{6}}3{{s}^{1}}\\N{{a}^{+}}=1{{s}^{2}}2{{s}^{2}}2{{p}^{6}}\end{array}\)

Neutral Na atom has the tendency to acquire the stable s2p6 configuration by the loss of one electron. Due to s2p6 configuration of Na+, the further separation of electron is difficult. So IAS group metals form mono positive ions.

(b)   If the difference of two consecutive I.P.s. of an element is 11.0 eV or less, the higher oxidation state is stable. For e.g. the difference of first and second I.P of Mg is 7.4 eV which is less than 11.0 eV. So Mg2+ will be stable. It can also be explained on the basis of its electronic configuration.

The electronic configuration of Mg2+ is stable s2p6 configuration Mg2+ = 1s22s22p6 So IIA group elements form di positive ions.

(c)   The difference of first and third I.P. of Al is 12.8 eV which is more than 11eV. Therefore first Oxidation state of Al i.e. Al+ must be stable in gaseous. This is due to the proportions distribution of lattice energy and the difference of second and third.

Electron affinity (EA) is the amount of energy released when a neutral isolated gaseous atom accepts an electron to form gaseous anion,

\(X(g)\,+e\,\to {{X}^{-}}(g)\,+EA\)

Similarly, second and third electron can be added to form gaseous dinegative and trinegative ions. The energy changes accompanying the addition of first, second, third, etc. electrons to neutral isolate gaseous atoms are called successive electron affinities and are designated as EA1, EA2 etc.

Since an atom has a natural tendency to accept an electron, therefore, the first electron affinity  (EA1) is always taken as positive. However, the addition of second electron to the negatively charged ion is opposed by coulombic repulsion and hence energy has to be supplied for the addition of second electron. Thus, second electron affinity (EA2) of an elements is taken as negative. For example,

\(O(g)+{{e}^{-}}\to {{O}^{-}}(g);\,E{{A}_{1}}=+141KJ\,mo{{l}^{{-1}}}\)

The factors which affect the electron affinity are :

(i) Atomic size, (ii) Nuclear charge and (iii) Electronic configuration.

The EA1 of most of the elements are positive i.e., energy is evolved i.e. \(\Delta H\) is negative. The EA2 of an element is always negative, i.e. energy is absorbed i.e.  is positive.

Periodic variation of electron affinity :

(i) In general, the electron affinity decreases in a group from top to bottom and increases along a period from left to right.

(ii) The electron affinity of element of second period are, however, lower than those of the corresponding elements of the third period due to small size as a result of which it has strong electron – electron repulsions, e.g. the electron affinity of F is lower than that of Cl The actual order being  Cl> F> Br> I.

(iii) The electron affinities of elements having exactly half – filled and completely filled orbitals are essentially zero, i.e. electron affinities of Be, Mg, Ca  etc.  N and inert gases are zero.

(iv) The electron affinities of elements of group 15 increase as we move down the group from P to Bi.

(6) Electronegativity (EN) : It is the tendency of an atom in a molecule to attract the shared pair of electrons towards itself.

The factors which affect the Electronegativity are :  (i) atomic size and (ii) nuclear charge.

Periodic variation of Electronegativity :

(i) The electronegativity of an atom decreases regularly down a group from top to bottom but increase along a period from left to right.

(ii) The electronegativity of an atom depends upon its bonding state, e.g. the electronegativity of sp-hybridized carbon is the maximum followed by sp² and then sp³.

Applications of electronegativity :

(i)  It helps to predict the polarity of bonds and dipole moment of molecules. If the EN difference between the two atoms is more than 1.7 the bond is considered to be ionic since it has more than 50% ionic character. If, however, the difference is less than 1.7, it is considered to be covalent.

(ii) It helps to have an idea about the atomic size. Higher the electronegativity, smaller is the atomic size.

(iii) Similarly, it helps to have an idea about the bond length. Higher the electronegativity difference, smaller is the bond  length.

       Electronegativity of an element can be measured by either of the following three methods :


Mulliken scale: On the Mulliken scale electronegativity \((\chi )\) is taken as average of IE and EA i.e. \(=\frac{{IE+EA}}{2}\)

  • where IE and EA are expressed in electron volts.
  • Pauling scale : This is the most widely used scale and is based upon bond energy data. According to Pauling, the difference in electronegativity of two atoms A and B is given by the relationship.

\({{\chi }_{B}}-{{\chi }_{{A=0.208\sqrt{\Delta }}}}\) Where  \(\Delta ={{E}_{{A\,-\,B}}}\sqrt{{{{E}_{{A\,-A\,}}}\times {{E}_{{B\,-\,B}}}}}\)

Here \({{E}_{{A\,-B,\,\,}}}{{E}_{{A-A}}}\) and \({{E}_{{B-B}}}\) represent bond dissociation energies of the bonds, A–B, AA, and BB respectively. By giving a reference value of 2.1 to H, the electrognegativities of almost all the elements have been calculated. On this scales fluorine has maximum value of 4.0.

If,                             a) \({{X}_{B}}-{{X}_{A}}=1.7,\,\,50\%\) ionic

b) \({{X}_{B}}-{{X}_{A}}>1.7,\) predominantly ionic

c) \({{X}_{B}}-{{X}_{A}}<\,\,1.7\) predominantly covalent

The Pauling and the Mulliken scales are related by the following expression : \({{\chi }_{{\text{Pauling}}}}\text{ }=\frac{{{{\chi }_{{\text{Mulliken}}}}}}{{2.8}}\)

(iii) Allred and Rochow’s method. According to this method, the electronegativity of an element A is given by the relation, \({{\chi }_{A}}=0.744+\frac{{0.359{{Z}_{{wff}}}}}{{{{r}^{2}}}}\) Where \({{Z}_{{eff}}}\) is the effective nuclear charge which is calculated with the help of Slater’s rules and ‘r’ is the covlent radius of the atom in angstroms.

Some other periodic properties

  • Atomic volume : It is defined as the volume occupied by one gram atom of an element. Mathematically,

Atomic volume = gram atomic wt./Density in solid state

Units of atomic volume are c.c./mole. Atomic volume signifies the volume occupied by one mole (Avogadro number) of atoms of the given element in solid state. Lower atomic volume generally leads to higher density, increased hardness and brittleness, higher melting and boiling points, less malleability and ductility.

(i) While descending a group, the atomic volume generally incre

ases which is due to increase in the number of shells though the valence electrons in a given group remains constant.

(ii) While going left to right across a period the atomic volume first decreases to a minimum and then increases. Francium has the highest atomic volume and boron has lowest atomic volume.

(2) Density :  The density of the elements in solid state varies periodically with their atomic numbers. At first, the density increases gradually in a period and becomes maximum somewhere for the central members and then starts decreasing afterwards gradually.

(3) Melting and boiling points : The melting points of the elements exhibit some periodicity with rise of atomic number. It is observed that elements with low values of atomic volumes have high melting points while elements with high values of atomic volumes have low melting points. In general, melting points of elements in any periodic at first increase and become maximum somewhere in the centre and thereafter begins to decreases.

Tungsten has the maximum melting point (3410°C) amongst metals and carbon has the maximum melting point (3727°C) amongst non-metals. Helium has the minimum melting point (–270°C). The metals,  and Hg are known in liquid state at 30°C.

The boiling points of the elements also show similar trends, however, the regularities are not so striking as noted in the case of melting points.

(4) Oxidation state (Oxidation number, O.N.) : Oxidation number of an element in a compound is the total number of electrons it appears to have gained or lost (negative and positive oxidation states respectively) during the formation of that particular compound.

(5) Magnetic properties : Magnetic properties of matter depend on the properties of the individual atoms. A substance (atom, ion or compound) capable of being attracted into a magnetic field is known as paramagnetic. The paramagnetic substances have a net magnetic moment which in turn is due to the presence of unpaired electron(s) in atoms, ions or molecules. Since most of the transition metal ions have unpaired d–electrons, they show paramagnetic behaviour. The exceptions are \(S{{c}^{{3+}}},\,T{{i}^{{4+}}},\,Z{{n}^{{2+}}},\,C{{u}^{+}},\) etc. which do not contain any unpaired electron and hence are diamagnetic.

On the other hand, a substance which is repelled by a magnetic field is known as diamagnetic. Such substances do not have any net magnetic moment because they do not have any unpaired electron. Electrons determine the magnetic properties of matter in two ways,

  • Each electron can be treated as a small sphere of negative charge spinning on its axis. The spinning of charge produces magnetic moment.
  • An electron travelling in closed path around a nucleus will also produce magnetic moment just as does electric current travelling in a loop of wire.

The observed magnetic moment is therefore the sum of the two moments: the spin moment and the orbital moment. It is expressed in units called Bohr Magnetons (BM). In terms of n (number of unpaired electron), magnetic moment is given by the formula, \(\mu =\sqrt{{n\,(n+2)}}\)

Greater the number of unpaired electrons in a substance, the greater is the magnetic moment of the substance. The value of magnetic moment has been used to calculate the number of unpaired electrons in an ion. In some cases, even the structure of the molecule or complex is indicated by its magnetic moment.

Paramagnetism is generally measured by a simple device known as Guoy’s balance which involves weighing the species in presence of a magnetic field.

Ferromagnetism is a special property observed in some substances in the solid state. Such substances are strongly attracted to magnetic field and may retain the magnetic properties for some time even after the removal of the field. The most common example is of Fe followed by Co and Ni.

(6) Hydration and hydration energy

(i) Hydration energy is the enthalpy change that accompanies the dissolving of 1 mol of gaseous ions in water.

(ii) Size of ions and its charge determines extent of hydration. Greater the charge smaller the size of the ion, greater the attraction for the lone pair of O of H2O, hence greater the extent of hydration energy.

(a) Size of the hydration ion increases.

(b) Ionic mobility decreases i.e. heavier (hydrated) ions moves slower.

(7) Acid-base-character of oxides

(i) On moving across a period, the basic character of the oxides gradually changes first into amphoteric and finally into acidic character.

(ii) On moving down a group, reverse behaviour is observed i.e., from more acidic to more basic.

(iii) Stability of oxides decreases across a period.

(8) Hydrides

(i) Hydrogen combines with a number of other elements including metals and non-metals to form compounds called hydrides.

(ii) Covalent nature of hydrides increases across a period and decreases down the group.

(iii) Ionic hydride are better reducing agents than covalent hydride and reducing nature of hydride decreases across a period and increases down the group.

Diagonal relationship

Certain elements of 2nd period show similarity with their diagonal elements in the 3rd period as shown below :

Thus, Li resembles Mg, Be resembles Al and B resembles Si. This is called diagonal relationship and is due to the reason that these pairs of element  have almost identical ionic radii and polarizing power (i.e. charge/size ratio). Element of second period are known as bridge elements.

Anomalous behaviour of the first elements of a group: The first element of a group differs considerably from its congeners (i.e. the rest of the element of its group). This is due to (i) small size (ii) high electronegativity and (iii) non availability of d-orbitals for bonding. Anomalous behaviour is observed among the second row elements (i.e. Li to F).

Question      “Electron affinity of Cl is the highest among the halogens – yet, F is the strongest oxidising agent” . Why?

Solution          The electron affinity of Cl is highest which is also the measure of oxidising ability. The strength of oxidising agent (i.e. its oxidation potential) depends on several energy terms and best represented by Born Haber type of cycle.





F has low heat of dissociation (159 kJ mol–1) and heat of hydration ( – 515 kJ mol1) and overall position is that largest  value in the energy cycle ( –726 kJ mol–1/Cl, 607 kJ mol–1)

Question  \(\displaystyle \frac{{{{N}_{0}}}}{2}\) atoms of (X)g are converted into (X+)g by energy  \(\displaystyle \Delta {{H}_{1}}\,\,\,\,and\,\frac{{{{N}_{0}}}}{2}\)  atoms of  (X+)g are converted into (X)g by energy \(\displaystyle \Delta {{H}_{2}}.\)

Calculate (IE) and (EA) of (X)g.

Solution          Let (IE) of (X)g = I per atom and (EA) of (X)g = – E per atom

i) \(\displaystyle {{\left( X \right)}_{g}}\xrightarrow{{}}{{\left( {{{X}^{+}}} \right)}_{g}}+{{e}^{+}}\) Energy required to ionize  \(\displaystyle \frac{{{{N}_{0}}}}{2}atoms\,\,{{\left( X \right)}_{g}}=\frac{{{{N}_{0}}I}}{2}\)

\(\displaystyle \frac{{{{N}_{0}}I}}{2}=\Delta {{H}_{1}}\)

\(\displaystyle I=\frac{{2\Delta {{H}_{1}}}}{{{{N}_{0}}}}\)

ii) \(\displaystyle {{X}_{{\left( g \right)}}}\xrightarrow{{}}X_{{\left( g \right)}}^{+}+{{e}^{-}}Energy=\frac{{{{N}_{0}}I}}{2}for\frac{{{{N}_{0}}}}{2}atoms\)

\(\displaystyle {{X}_{{\left( g \right)}}}+{{e}^{-}}\xrightarrow{{}}X_{{\left( g \right)}}^{-}Energy=-\frac{{{{N}_{0}}E}}{2}of\frac{{{{N}_{0}}}}{2}atoms\)

\(\displaystyle \frac{{{{N}_{0}}l}}{2}-\frac{{{{N}_{0}}E}}{2}=\Delta {{H}_{2}}\)

\(\displaystyle -\frac{{{{N}_{0}}E}}{2}=\Delta {{H}_{2}}-\Delta {{H}_{1}}\)

\(\displaystyle -E=\frac{{2\left( {\Delta {{H}_{2}}-\Delta {{H}_{1}}} \right)}}{{{{N}_{0}}}}\)

(IE) ionisation energy= \(\displaystyle \frac{{2\Delta {{H}_{1}}}}{{{{N}_{0}}}}ato{{m}^{{-1}}}\)

(EA) electron affinity = \(\displaystyle \frac{{2\left( {\Delta {{H}_{2}}-\Delta {{H}_{1}}} \right)}}{{{{N}_{0}}}}ato{{m}^{{-1}}}\)

Question        The sums of first and second ionisation energies and those of third fourth ionisation energies in (MJ mol–1) of nickel and platinum are

                        (IE)1 + (IE2)                 (IE3) + (IE4)4

Ni 2.49                         8.80

Pt 6.70                          2.66

Based on this information; write

(i)  The most common oxidation states of Ni and Pt.

(ii) Names of metal (Ni or Pt) which can more easily form compounds in its +4oxidation state.

Solution          (i)  Ni = +2, Pt =+4 since (IE)1 + (IE)2 of Ni is less than its (IE)3 + (IE)4 and re verse is the case in pt.

(ii) Platinum form more stable complex in +4 states due to its higher stability than +2  states.

Question        In alkali metal group which is the strongest reducing agent and why?

Solution          Li is the strongest reducing agent. Since I.P. decreases down the group we would expect that Li will have the lowest reducing power in that group. But group, Li will have highest reducing power.

Question        The electron affinity of sulfur is greater than oxygen. Why?

Solution          This is because of smaller size of oxygen due to which it has go t higher change density and there is electronic repulsion as it takes electron So its E.A. is less than sulphur.