Electronegativity

Here, you will learn about…

• What is electronegativity? – Definition with Example
• Measurement of electronegativity
  • Pauling’s Scale of Electronegativity
  • Mulliken Scale of Electronegativity
  • Allred and Rochow’s Scale of Electronegativity
• Factors upon which electronegativity depends
• Variation of electronegativity in the period table
• Applications of the Electronegativity
• FAQs

What is Electronegativity? – Definition

The electronegativity of an element is defined as a measure of the tendency of an atom in a molecule to attract the shared electron pair towards itself when it is covalently bonded with another atom.

Thus, in a heteronuclear bond A – B between two different atoms A and B, the shared electron pair will be shifted more toward the element with greater electronegativity.

For example, if atom ‘A’ has a stronger tendency to attract the shared electron pair towards itself as compared to that of atom B, then the electronegativity of atom ‘A’ is greater than atom ‘B’. This leads to an increase in the partial negative charge on atom ‘A’ and a decrease in the partial negative charge on atom ‘B’ or an increase in the partial positive charge on atom ‘B’. It can be represented as .

The electronegativity of an element is not an inherent property and it depends on the bonding environment of that element. For example, the electronegativity of the ‘P‘ atom in the compounds PCl₃ and PCl₅ are not the same.

Measurement of Electronegativity

The absolute magnitude of the electronegativity value of an element is however not fixed and it depends on the bonding environment of that element. There are many theoretical approaches that have been proposed to calculate electronegativity. These are well-known as different electronegativity scales.

Pauling’s Scale of Electronegativity

In 1932, Pauling first introduced an electronegativity scale based on the observed bond dissociation energy. He first proposed that the bond dissociation energy of the purely covalent bond in a molecule A – B is the geometric mean of the bond dissociation energies of the purely covalent molecules A – A and B – B, . This is only true when two 100% covalent molecules, A – A and B – B react together and form a 100% covalent molecule A-B (A-A + B-B → A – B + A – B). But, Pauling noticed that when molecules A-B is not purely covalent, the bond dissociation energy of A – B is always greater than that of the geometric mean of the bond dissociation energies of the purely covalent molecules A– A and B – B.

This indicates that the A-B bond is more stable when it is not purely covalent. It is also well known that ionic bonds are more stable than covalent bonds. Therefore, the extra stability in the A-B was derived from the unequal sharing of the bonding electrons between two dissimilar elements A and B, giving rise to the ionic character in the A-B bond.

Thus, the difference in bond dissociation energy between A-B and the geometric mean of A – A and B – B bonds increases when the ionic character of the A-B bond increases. This difference in energy due to the partial ionic character of the A-B is called ionic resonance energy (∆). The ionic resonance energy (∆) can be represented as

—(1)

Where,
= The actual bond dissociation energy of the A-B bond.

= Geometric mean of the bond dissociation energies of the purely covalent molecules A – A and B – B.

The contribution of ionic character in the A-B bond is due to the electronegativity difference between A () and B (). Thus, ∆ is represented as the measure of the difference in electronegativity of the bonded atoms A and B. From this consideration, Pauling proposed an empirical relation:

—-(2)

From equations (1) and (2),

Where, ∆ is expressed in kcal mol ^{– 1}.

If the ∆ is expressed in kJ/mole, the above equation can be expressed as follows.

(1 eV = 96.48 kJ/mole = 23.06 kcal/mole)

Since from the above equation, it is only possible to determine the electronegativity difference, the electronegativities of all elements can be determined by taking the electronegativity value of at least one element as a reference point. Pauling assigned the electronegativity value of F atom 4.0 and calculated the electronegativities of all other elements. The electronegativities of different elements in the periodic table based on Pauling’s electronegativity Scale are represented in the following table.

Atomic Number	Name of the Atom (Symbol)	Electronegativity Value
1	Hydrogen (H)	2.2
2	Helium (He)	–
3	Lithium (Li)	0.98
4	Beryllium (Be)	1.57
5	Boron (B)	2.04
6	Carbon (C)	2.55
7	Nitrogen (N)	3.04
8	Oxygen (O)	3.44
9	Fluorine (F)	3.98
10	Neon (Ne)	–
11	Sodium (Na)	0.93
12	Magnesium (Mg)	1.31
13	Aluminium (Al)	1.61
14	Silicon (Si)	1.9
15	Phosphorus (P)	2.19
16	Sulphur (S)	2.58
17	Chlorine (Cl)	3.16
18	Argon (Ar)	–
19	Potassium (K)	0.82
20	Calcium (Ca)	1
21	Scandium (Sc)	1.36
22	Titanium (Ti)	1.54
23	Vanadium (V)	1.63
24	Chromium (Cr)	1.66
25	Manganese (Mn)	1.55
26	Iron (Fe)	1.83
27	Cobalt (Co)	1.88
28	Nickel (Ni)	1.91
29	Copper (Cu)	1.9
30	Zinc (Zn)	1.65
31	Gallium (Ga)	1.81
32	Germanium (Ge)	2.01
33	Arsenic (As)	2.18
34	Selenium (Se)	2.55
35	Bromine (Br)	2.96
36	Krypton (Kr)	3
37	Rubidium (Rb)	0.82
38	Strontium (Sr)	0.95
39	Yttrium (Y)	1.22
40	Zirconium (Zr)	1.33
41	Niobium (Nb)	1.6
42	Molybdenum (Mo)	2.16
43	Technetium (Tc)	1.9
44	Ruthenium (Ru)	2.2
45	Rhodium (Rh)	2.28
46	Palladium (Pd)	2.2
47	Silver (Ag)	1.93
48	Cadmium (Cd)	1.69
49	Indium (In)	1.78
50	Tin (Sn)	1.96
51	Antimony (Sb)	2.05
52	Tellurium (Te)	2.1
53	Iodine (I)	2.66
54	Xenon (Xe)	2.6
55	Cesium (Cs)	0.79
56	Barium (Ba)	0.89
57	Lanthanum (La)	1.1
58	Cerium (Ce)	1.12
59	Praseodymium (Pr)	1.13
60	Neodymium (Nd)	1.14
61	Promethium (Pm)	1.13
62	Samarium (Sm)	1.17
63	Europium (Eu)	1.2
64	Gadolinium (Gd)	1.2
65	Terbium (Tb)	1.22
66	Dysprosium (Dy)	1.23
67	Holmium (Ho)	1.24
68	Erbium (Er)	1.24
69	Thulium (Tm)	1.25
70	Ytterbium (Yb)	1.1
71	Lutetium (Lu)	1.27
72	Hafnium (Hf)	1.3
73	Tantalum (Ta)	1.5
74	Tungsten (W)	2.36
75	Rhenium (Re)	1.9
76	Osmium (Os)	2.2
77	Iridium (Ir)	2.2
78	Platinum (Pt)	2.28
79	Gold (Au)	2.54
80	Mercury (Hg)	2
81	Thallium (Tl)	1.62
82	Lead (Pb)	2.33
83	Bismuth (Bi)	2.02
84	Polonium (Po)	2
85	Astatine (At)	2.2
86	Radon (Rn)	–
87	Francium (Fr)	0.7
88	Radium (Ra)	0.89
89	Actinium (Ac)	1.1
90	Thorium (Th)	1.3
91	Protactinium (Pa)	1.5
92	Uranium (U)	1.38
93	Neptunium (Np)	1.36
94	Plutonium (Pu)	1.28
95	Americium (Am)	1.3
96	Curium (Cm)	1.3
97	Berkelium (Bk)	1.3
98	Californium (Cf)	1.3
99	Einsteinium (Es)	1.3
100	Fermium (Fm)	1.3
101	Mendelevium (Md)	1.3
102	Nobelium (No)	1.3
103	Lawrencium (Lr)	–
104	Rutherfordium (Rf)	–
105	Dubnium (Db)	–
106	Seaborgium (Sg)	–
107	Bohrium (Bh)	–
108	Hassium (Hs)	–
109	Meitnerium (Mt)	–
110	Darmstadtium (Ds)	–
111	Roentgenium (Rg)	–
112	Copernicium (Cn)	–
113	Nihonium (Nh)	–
114	Flerovium (Fl)	–
115	Moscovium (Mc)	–
116	Livermorium (Lv)	–
117	Tennessine (Ts)	–
118	Oganesson (Og)	–

Mulliken Scale of Electronegativity

In 1934, R.S. Mulliken proposed a new scale for electronegativity based on the ionization energy and electron affinity of an element. According to the Mulliken scale,

Where I.E. = Electron Affinity in eV per atom and E.A. = Electron Affinity in eV in per atom.

If I.E. and E.A. are expressed in kJ. mol^-1, then

1 eV per molecule = 96.48 kJ. mol^-1.

Similarly, If I.E. and E.A. are expressed in kcal. mol^-1, then

1 eV per molecule = 23.06 kcal. mol^-1.

Allred and Rochow’s Scale of Electronegativity

In 1958, Allred and Rochow introduced a new electronegativity scale based on Coulomb’s law of attraction. According to them, the attraction of bonding electrons of an atom can be expressed based on Coulomb’s law of attraction.

According to Coulomb’s law,

—(1)

Where,

Z_eff is the effective nuclear charge which can be calculated from the Slater rules. Z_eff = Z (Nuclear charge) – σ (Screening Constant).

e = Charge of the electron

r = Radius of the atom

Based on equation 1, the following mathematical equation is introduced by Allred and Rochow:

Factors on which electronegativity depends

1. Valency state and oxidation state

The electronegativity depends on the oxidation state or valency state of an element.

The elements with a higher oxidation state have a greater tendency to attract the electron towards themselves. Therefore, an element with a higher oxidation state is more electronegative than that of an element with a lower oxidation state.

For example, phosphorus (P) in PCl₅ is more electronegative than phosphorus (P) in PCl₃. The oxidation state of P in PC₅ is +5 and in PCl₃ is +3.

2. Size of the atoms

The electronegativity decreases with the increase in the size of the atom and vice–versa. Since small atoms attract electrons more strongly than large ones and are therefore more electronegativity.

3. Nature of the hybridization

The electronegativity of an element also depends on the nature of the hybridization of bonding atoms. When % of s character in hybrid orbital increases, the distance between two atoms decreases, and electron attracting power increases. Thus, the higher the % of s character of the hybrid orbitals higher will be its electronegativity.

For example, let us consider the three simplest hydrocarbons, Ethane (C₂H₆), Ethene (C₂H₄), and Ethyne (C₂H₂) where hybridization of C atom sp³, sp², sp in Ethane, Ethene, Ethyne respectively. The % of s characters in sp, sp², and sp³ hybridized orbitals are 50%, 33.33%, and 25%, respectively. Therefore, the electronegativity order of the C atom is Ethyne (sp) > Ethene (sp²)> Ethane (sp³).

Variation of electronegativity in the period table

In a period

In the periodic table, the electronegativity gradually increases along the period moving from left to right. This is due to the increase in effective nuclear charge and decrease in the size of the atoms from left to right in a period. Thus, alkalies in the left are the least electronegative elements and halogens in the right are the most electronegative elements.

In a group

In moving down in a group, the electronegativity gradually decreases. This is due to the increase in the size of the atoms moving down in a group. Since small atoms attract electrons more strongly than large ones. Therefore, a lower element of a group is less electronegative than the upper element.

The periodic trend of electronegativity is summarized in the following diagram.

Applications of the Electronegativity

Electronegativity values are very helpful in understanding and predicting various properties related to the energy and charge distribution and chemical bonds, e.g. bond polarity, bond dissociation energy, and bond moment.

FAQs

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