Electronegativity

Electronegativity is a qualitative chemical property that measures the relative tendency of an atom in a chemical compound to attract the shared pair of electrons toward itself in a covalent bond. Unlike electron affinity, electronegativity is not an inherent property of an isolated atom; it is a context-dependent property exhibited by an atom only when it is chemically bonded to another element.

Scales for Measuring Electronegativity

Because electronegativity is a relative, dimensionless property, it cannot be measured directly in energy units. Scientists have developed several empirical scales to quantify it.

Pauling Scale

Developed by Linus Pauling in 1932, this is the most widely used scale. It is based on the excess bond dissociation energies (Δ) of heteronuclear diatomic molecules compared to their constituent homonuclear molecules. The electronegativity difference between two atoms A and B is calculated using the formula:

Where Δ (in kJ/mol) represents the resonance energy:

• Reference Standard: Pauling arbitrarily assigned the highest value of $4.0$ to Fluorine (F), the most electronegative element, to calibrate all other elements.

Mulliken-Jaffe Scale

Proposed by Robert Mulliken, this scale treats electronegativity as an inherent atomic property by averaging the ionization enthalpy (IE) and electron gain enthalpy (Δ H_eg) of an isolated atom.

• Conversion: Mulliken values are roughly $2.8$ times larger than Pauling values. The conversion formula is:

Allred-Rochow Scale

This scale defines electronegativity as the electrostatic force exerted by the effective nuclear charge (Z_eff) of an atom on the valence electrons shared in a covalent bond.

Where r is the covalent radius of the atom in Angstroms (Å).

Key Factors Influencing Electronegativity

Effective Nuclear Charge (Z_eff)

Electronegativity is directly proportional to the effective nuclear charge. A higher positive charge in the nucleus exerts a stronger electrostatic pull on the bonding pair of electrons.

Atomic Radius

Electronegativity is inversely proportional to atomic size. In smaller atoms, the shared electron pair is located closer to the positively charged nucleus, resulting in a stronger force of attraction.

Oxidation State

The electronegativity of an element increases as its positive oxidation state increases. For instance, an Fe³⁺ ion is more electronegative than an Fe²⁺ ion because the higher positive charge increases the atom’s pull on external electrons.

Hybridization State

The electronegativity of an atom varies with the type of hybridization, depending on the percentage of s-character. Because s-orbitals are closer to the nucleus than p-orbitals, a higher s-character increases electronegativity.

Periodic Trends in Electronegativity

Variations Across a Period (Left to Right)

• Trend: Electronegativity increases significantly.
• Reason: Atomic radius decreases and the effective nuclear charge increases. Consequently, the nucleus can attract the shared electron pair with greater force.

Variations Down a Group (Top to Bottom)

• Trend: Electronegativity decreases.
• Reason: Atomic radius increases due to the addition of new electronic shells, and the shielding effect increases. The nucleus is further removed from the shared valence electrons, reducing its pulling force.

Exceptions and Anomalies

• Noble Gases: Elements of Group 18 (He, Ne, Ar) do not have standard Pauling electronegativity values assigned to them because they form very few stable chemical bonds.
• Transition Elements: Electronegativity changes minimally across transition series (d-block) due to the counterbalancing effect of inner d-electrons shielding the outer electrons.
• Post-Transition Metals: Elements like Gallium (Ga) have a slightly higher electronegativity than Aluminum (Al) due to the poor shielding effect of $3delectrons, which causes a contraction in atomic size (d-block contraction). </li> </ul> <h4>Applications of Electronegativity in Chemistry</h4> <h4>Prediction of Bond Nature</h4> <p> The difference in electronegativity (\Delta \chi) between two bonded atoms determines the ionic or covalent character of the chemical bond. </p> <table> <thead> <tr> <td><strong>Electronegativity Difference (Δχ)</strong></td> <td><strong>Nature of the Bond</strong></td> <td><strong>Example</strong></td> </tr> </thead> <tbody> <tr> <td><b>\Delta \chi = 0</b></td> <td>Purely Covalent / Non-polar</td> <td>\text{H}_2, \text{Cl}_2</td> </tr> <tr> <td><b>\Delta \chi < 1.7</b></td> <td>Polar Covalent</td> <td>\text{H}_2\text{O}, \text{HCl}</td> </tr> <tr> <td><b>\Delta \chi = 1.7</b></td> <td>50\%Ionic and50\%Covalent</td> <td>\text{HF}</td> </tr> <tr> <td><b>\Delta \chi > 1.7</b></td> <td>Predominantly Ionic</td> <td>\text{NaCl}, \text{CsF}</td> </tr> </tbody> </table> <h4>Metallic and Non-Metallic Character</h4> <ul> <li> Elements with very low electronegativity values easily lose electrons and are classified as <b>metals</b> (electropositive elements). </li> <li> Elements with high electronegativity values readily gain or attract electrons and are classified as <b>non-metals</b>. </li> </ul> <h4>Bond Length and Bond Strength</h4> <p> A larger electronegativity difference between two bonding atoms increases the electrostatic attraction between them. This draws the atoms closer together, shortening the bond length and increasing the overall bond strength. </p> <h4>UPSC Prelims Fact File and Trivia</h4> <h5>Extremes of the Periodic Table</h5> <ul> <li> <b>Most Electronegativity Element:</b> Fluorine (\text{F}) has the highest value of %%MONEY_BLOCK₂%% on the Pauling scale, followed by Oxygen (%%MONEY_BLOCK₃%%) and Nitrogen (%%MONEY_BLOCK₄%%). </li> <li> <b>Least Electronegative Element:</b> Cesium (\text{Cs}) and Francium (\text{Fr}$) hold the lowest electronegativity values ($0.7$ on the Pauling scale), making them the most electropositive elements.

Top Electronegativity Order (FONClBrISCH)

UPSC aspirants should note the decreasing order of electronegativity for the most common non-metals:

Diagonal Relationships

Certain second-period elements show similar electronegativity values to third-period elements placed diagonally opposite to them due to a balancing of size and charge effects (e.g., Lithium and Magnesium; Beryllium and Aluminum).

Last Modified: May 25, 2026