B.sc 1st year Book (Page 8)
Electronegativity
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.
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
$$ X_A-X_B\:\infty \:\sqrt{\Delta _{A-B}}\:Or\:X_A-X_B=\sqrt{\Delta \:_{A-B}}……\left(iv\right) $$
or
$$ X_A-X_B=K\sqrt{E\:_{A-B}\sqrt{E_{A-A}\times E_{B-B}}}……\left(v\right) $$
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$$ X_A-X_B=0.208\sqrt{\Delta \:_{A-B}}……\left(vi\right) $$
Or
$$ X_A-X_B=0.208\sqrt{E\:_{A-B}\sqrt{E_{A-A}\times E_{B-B}}}……\left(vii\right) $$
Or
$$ X_A-X_B=0.182\sqrt{E\:_{A-B}\sqrt{E_{A-A}\times E_{B-B}}}……\left(viii\right) $$
This Pauling scale gives the value of the difference between electronegativities of two atoms ‘A’ and ‘B’. 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 teams 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) |
| I A | ||||||
| H 2.1 | II A | III A | IV A | V A | VI A | VII A |
| 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 :
. 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
(a) For a pure covalent bond i.e.
The disadvantage of Mulliken’s 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):
Variation of ionization energy values :
(a) ln a period :
On proceeding from left to right, the IE values in general increase with a gradual increase in the magnitude of nuclear charge. Consequently, If we consider the IE values of the elements of the second period, these should increase in the following order:
| IA | IIIA | IIA | IVA | VIA | VA | VIIA | Zero | |
| 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
(ii) Charge on the atom :
(iii) Hybridization :
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.
(iv) Color of salts :
(v) Diagonal relationship :
(vi) Metallic character of the elements:
(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 :
- 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+decreases in the following order :
- 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.
- 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.
- 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.
- 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
- 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.
- 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”.
| 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 |
| 9 |