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Let us consider the formation of a molecule by sharing of electrons between two atoms. This sharing of electrons may take place between either two similar atoms or dissimilar atoms. In the case of two similar atoms, the shared pair of electrons is equally attracted by both atoms and lies at the center of the molecule e.g. H2 molecule. But in the case of molecules formed by dissimilar atoms, the shared pair of electrons is nearer to one of them which has a greater attracting tendency viz. In HCl, it lies nearer to the Cl atom. In other words, chlorine is more electronegative than hydrogen atoms.

Thus, electronegativity is defined as the tendency (or ability) of an atom to attract a shared pair of electrons towards itself in a molecule’, In general.

Electronegativity, Electronegativity structure,

The electronegativity of an atom ‘ A ‘ is generally represented as XA where ‘ A ‘ is an atom. Since it is the tendency of an atom hence it has no unit of measurement.

The scale of Electronegativity (EN):

The EN scales have been given with the help of various types of experimental data. which are completely different.

(1) Pauling’s bond energy scale :

In 1932 Pauling defined electronegativity’s ‘a power of a form in its molecule to attract bonding electrons pair towards itself. He devised an electronegativity scale on the basis of bond energy. According to him, if a molecule A−B is formed by the combination of A2 and B2 molecules of which EA−B is the bond energy of the AB molecule, EA−A and EB−B is the bond energies of A2 and B2 molecules respectively then the bond energy of the molecules AB is a geometric mean of the bond energies of the A2 and B2 molecules.

$$ i.e\:E_{A-B}=\sqrt{E_{A-A}}\times \sqrt{E_{B-B}}……..\:\left(i\right) $$

But experimentally, it is found that EA-B is greater than the calculated value.

$$ i.e\:E_{A-B}>\sqrt{E_{A-A}}\times \sqrt{E_{B-B}}……..\:\left(ii\right) $$

The difference between EA-B and √EA-A x EB-B is called the ‘ionic resonance energy’ of the A-B, bond and is represented as DA-B. Thus

$$ \Delta _{A-B}=E_{A-B}-\sqrt{E_{A-A}}\times \sqrt{E_{B-B}} $$

= 23(xA – xB)2……… (iii)

Pauling further suggested that the square root DA-B i.e. √ΔA-B is the measure of partial ionic character. As the electronegativity difference of the atoms A and B increase the partial ionic character of the A—B bond also increases. Let us consider that XA and XB are electronegativities of atoms A and B respectively. then

$$ X_A-X_B\:\infty \:\sqrt{\Delta _{A-B}}\:Or\:X_A-X_B=\sqrt{\Delta \:_{A-B}}……\left(iv\right) $$


$$ X_A-X_B=K\sqrt{E\:_{A-B}\sqrt{E_{A-A}\times E_{B-B}}}……\left(v\right) $$

ΔA-B does not possess additive property i.e If we consider three covalent bonds via A—B, C—B, and C—A then

ΔA-B + ΔB-C ≠ ΔC-A : when XA > XB > XC

Where ‘K’ is proportionality constant and has a value equal to 0.208 If the bond energies are expressed in eV. The value of ‘K’ has been found to be equal to 0.182 when the bond energies are expressed in Kcal/Mole.

i.e$$ X_A-X_B=0.208\sqrt{\Delta \:_{A-B}}……\left(vi\right) $$


$$ X_A-X_B=0.208\sqrt{E\:_{A-B}\sqrt{E_{A-A}\times E_{B-B}}}……\left(vii\right) $$


$$ X_A-X_B=0.182\sqrt{E\:_{A-B}\sqrt{E_{A-A}\times E_{B-B}}}……\left(viii\right) $$

This Pauling scale gives the value of the difference between electronegativities of two atoms ‘A’ and ‘B’. While exploring such intricate concepts, some students might consider seeking professional assistance, perhaps even opting to have their masterarbeit schreiben lassen or master’s thesis written by experts, especially when dealing with complex scientific topics. We can Calculate the value of XA, If the value of XB is known. Pauling assigned an arbitrary value of electronegativity of fluorine which is equal to 4.1. He expressed the values of ionic resonance energies in terms of an electron volt (eV).

Where (1 eV = 23 Kcal per gram bond). The values of electronegativity of normal elements are given in table 2.12.

Table 2.12: Pauling scale of Electronegativity values of the normal Elements (Pauling Scale F = 4.1)
Li 1.00 Be 1.5 B 2.0 C 2.5 N 3.1 O 3.5 F 4.1
Na 0.97 Mg 1.2 Al 1.5 Si 1.7 P 2.1 S 2.4 Cl 2.8
K 0.90 Ca 1.0 Ga 1.6 Ge 2.0 As 2.2 Se 2.5 Br 2.7
Rb 0.89 Sr 1.0 In 1.5 Sn 1.72 Sb 1.82 Te 2.0 I 2.2
Cs 0.86 Ba 0.97 Ti 1.4 Pb 1.5 Bi 1.7 Po 1.8 At 1.8

(2) Mulliken’s Scale :

In 1934 Mulliken suggested an alternate approach to electronegativity. According to him, electronegativity is the mean of the difference between ionization energy and electron affinity of an atom. Let us consider two atoms ‘A’ and ‘B’ forming a molecule A−B, in which an electron is transferred from A to B resulting in the formation of A+cation and B−anion.
If IA is the ionization energy of atomA and (EA )Bis the electron affinity of atom Bthen, the change in energy is equal toIA−(EA)B-Alternatively, if the electron is transferred from B to A then IB−(EA)Ais the change in energy. Depending upon the ease of formation of these ions, eitherIA−(EA)B>IB−(EA)A or lB −(EA)A>IA−(EA)B. Mulliken further suggested that the electronegativity of an atom could be regarded as the average of ionization energy and electron affinity of an atom If xA and xB are the electronegativities of atoms A and B respectively. Then

xA=1A+(EA)A2 and xB=1B+(EA)B2

When the value of ‘ I ‘ and ‘EA’ are expressed in eV. It is found that the value of electronegativity of an atom measured by Mulliken is 2.8 times larger than the Pauling values hence he modified his scale as;
xA=1A+(EA)A2 and xB=1B+(EA)B2
If these values are measured in Kcal per mole then the above relation is expressed as:
xA=1A+(EA)A2 and xB=1B+(EA)B2
The conditions for the formation of purely covalent or ionic bonds on the basis of Mulliken’s scale are :
(a) For a pure covalent bond i.e. xA=xB
For ionic bond

The disadvantage of Mulliken’s Scale:

There are two main practical disadvantages of this scale:
(i) The value of electron affinities is not easily available
(ii) The value of IE and EA with reference to the transfer of electrons between the atomic orbitals is not always known due to the lack of their constitution.

(3) The Allred-Rochow electrostatic approach (1958):

Allred and Rochow’s scale is based on the electrostatic force between the nucleus and the valence electron. Thus, according to Allred and Rochow’s approach to the electrostatic force of attraction, FES (in dynes) between an atom ‘A’ and a bonding electron separated from its nucleus by its covalent radius {(rA)cov } in A, is the measure of electronegativity of the atom ‘A’. In academic settings, where understanding such complex concepts is crucial, many students often seek assistance, sometimes even opting to hausarbeit schreiben lassen for a deeper comprehension of such topics. Hence, according to Coulomb’s law –
For ionic bondThe Allred-Rochow electrostatic approach (1958)
FES values calculated from equation (ii) when plotted against the corresponding electronegativity values on Pauling scale, a straight line was obtained. From the respective slope and intercept of this straight line, we get
(XA)A-R = a . FES + b ………… (iii)

Variation of ionization energy values :

(a) ln a period :

On proceeding from left to right, the IE (Ionization Energy) values generally increase due to a gradual increase in the magnitude of nuclear charge. This trend is particularly notable when analyzing the elements of the second period. For students delving deeper into this topic, options like enlisting assistance with a seminararbeit schreiben lassen service can be beneficial, especially when compiling comprehensive research papers. Therefore, if we consider the IE values of the elements of the second period, these should increase in the following order:

Similarly, in the third period, I.E. values of the elements of the third period should increase in the following order:
Na < Be < B < C < N < O < F < Ne
However, because of the different factors which influence the values of IE, the actual order of IE in the 2nd and 3rd periods are:
IInd Period  Li < B < Be < C < O < N < F < Ne
I.E (In eV) 5.4 8.3 9.3 11.3 13.6 14.5 17.4 21.6
IIIrd Period Na < Al < Mg < Si < S < P < Cl < Ar
I.E (In eV) 5.1 6.0 7.6 8.1 10.4 11.0 13.0 15.8

Factors affecting the magnitude of electronegativity :

(i) Size of the atom:

In general, electronegativity decreases with an increase in the size of the atom, this is due to an increase in effective nuclear charge, thus

Size of the atom

(ii) Charge on the atom :

An atom while acquiring a positive charge, (either partial or integral), would tend to attract electrons more strongly than its parent ghostwriter jura atom. Thus, the greater the positive oxidation state of an atom greater would be its electronegativity. For example: In the oxides, NO2NO2 and N2O5 the oxidation state of nitrogen atoms is +2,+4, and +5 respectively. Hence, the nitrogen atom in N2O5 will be more electronegative:

(iii) Hybridization :

We know that s-electrons are more penetrating than the p-electrons therefore the magnitude of electronegativity. For example in carbon compounds, the central carbon atom may exhibit sp, sp2, or sp3 hybridization in which the contribution of scharacter is 50%, 33.3%, and 25% respectively. Thus, the electronegativity of carbon atoms increases with an increase in the multiplicity of the C-C bond. i.e.
Due to the greater electronegativity of carbon atoms in ethylene and acetylene, the electron pair involving the C−H bond is pulled more towards the carbon atom thereby C−H bond breaks more easily to release. H-atom as H+ ion Thus, CH4 is neutral, CH2=CH2 is slightly acidic and HC≡CH is considerably more acidic.
A similar, type of hybridization also affects the basicity of amines. The higher the s-character of hybrid orbitals of a nitrogen atom, the greater would be the electronegativity. Consequently, the lower would be the electron donation power of the nitrogen atom, and hence lower would be the basicity of the amine. For example, methyl amine (H3C−NH2) is more basic than methyl cyanide; H3C−C≡N, It is because CH3NH2 and H3C−C≡N involve sp3 and sp hybridization of the orbitals of nitrogen atoms respectively.
The electronegativity of an atom depends considerably upon the number and nature of atoms or groups to which it is attached. Thus, electronegativity is not a constant value For example, the electronegativity of phosphorus is 2.24 and 2.44 in PCl3 and PF3 molecules respectively. It is because the phosphorus atom in PF3 acquires a greater positive charge than in PCl3. While its accepted value is 2.25.
xA - xB values and % ionic character of molecule
It is concluded from the above table that when (xA−xB) = 1.7 the bond A−B is approximately 50% ionic and +50% covalent. When (xA−xB) is less than 1.7, the bond A−B will have more than 50% covalent character. On the other hand. when (xA−xB) is greater than 1.7, the A-B bond will be more than 50% ionic.

(iv) Color of salts :

The color of the salts also depends on upon electronegativity difference of the elements. If the electronegativity difference is more, the salt will be either light in color or colorless i.e. with an increase of covalent character the color of the ionic salt also changes from white to dark. For example, AgCl is white, AgBr is pale yellow and Agl is yellow it is because the electronegativity difference in Ag and Cl is more so it will be less covalent and more ionic. whereas the electronegativity difference in AgBr and Agl is less hence these are less ionic and more covalent thus it is concluded that pure ionic substances are colorless (white) and the cells have more covalent characters are colored. Transition metal salts and complexes are colored due to their more covalent nature.

(v) Diagonal relationship :

The elements diagonally placed in the 2nd and 3rd periods of the periodic table show many similar properties which is due to the almost the same electronegativity of the elements.

(vi) Metallic character of the elements:

The elements which possess less electronegativity are more electropositive metallic elements. While higher values of electronegativity are non-metals. Hence, the metallic character increases from top to bottom in a group and decreases on moving along a period from left to right.

(vii) Nomenclature of binary compounds :

Generally, binary compounds are regarded as derivatives of more electropositive elements: For example, the binary compounds of iodine and chlorine are represented as 1Cl and not Cl1 hence it is named iodine chloride. Similarly, oxygen difluoride is represented as OF2 and not as F2O.

(viii) Nature of XOH in aqueous solution :

The nature of the XOH molecule can be predicted on the basis of electronegativities of the atom X as below:

(a) If xO−xH > xX, the O-H bond will be more polar than the O-X bond hence ionization of XO−and H+ions will be obtained and the molecule will be acidic in nature.
(b) If xO−x4 < xO−xX, the O−H bond will be less polar than the O−X bond hence ionization of X+ and OH ions will be obtained and the molecule will be basic in nature.
Points to remember :

  1. A cation is always smaller than its corresponding atom. It is because a cation formed by the loss of electrons may result in the complete disappearance of the outermost shell. Therefore, the remaining inner shells do not extend so far in space which is why the cation becomes smaller than its corresponding atom.

Also, whenever a cation has formed the ratio of the nuclear charge to the number of electrons (z/e ratio) is increased which results from the effective nuclear charge is also increased and the electrons are pulled strongly towards the nucleus. Consequently, the cation becomes smaller:

2. An anion always targets its corresponding atom: because when an anion is formed by the addition of one or more electrons, the effective nuclear charge decreases, and the electron cloud expands which results from the anionic size increasing.

3. For isoelectronic fons (Ions having an equal number of electrons but different actual nuclear charge): the greater the nuclear charge, the greater will be the attraction for electrons and the smaller the ionic radius. For example, The ionic radius for the isoelectronic ions; C4., N3 .O2.F. Na2, Mg2+, and A2+ C4, N3, O2, F, Na2, Mg2+, and A2+decreases in the following order :


  1. The second ionization energy, f, is always more than the first, 11. Because in the second ionization potential, an electron is to be removed from a Unipositive ion which has more attraction for electrons, and thus the removal of this electron requires more energy.
  2. The ionization potentials of inert gases are very high: because these gases have energetically more stable electronic configurations and thus it becomes very difficult to remove an electron from them.
  3. Fluorine has a slightly lower electron affinity than chlorine: because fluorine has a very small atomic size and great electron density hence the incoming electron experiences greater repulsive forces due to electron-electron repulsion from the already present electrons than in chlorine.
  4.  Elements with nearly same electronegativity in the Pauling scale are :

  N=Cl=30;      C=S=1=2.5:     H=P=2.1:     Cs=Fr=0.7

  1. The first ionization potential of boron is less than beryllium: because in boron (1s2,2s′,2p′), the electron is to be removed from 2p which is very easy while in beryllium ( 182, 25″) electron is 50 be removed from 2 s which is difficult. In addition, Be has a more stable configuration than B.
  2. Amongst all elements of the periodic table –

(I)   helium has the smallest size
(II)  helium has the highest value of first ionization energy
(III) fluorine has the highest value of electronegativity
(IV) chlorine has the highest value of Irst electron affinity
(V)   fluorine is the most powerful oxidizing agent

10. Platinum is the most precious metal commonly known as “white gold”.

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