Electronegativity is defined as the relative tendency of an atom in a molecule to attract the bonded pair of electrons towards itself. For the electronegativity difference/polarity phenomenon to occur, the bonded atoms must be different. 

In case of a homonuclear diatomic molecule, e.g., H2, the bonded pair of electrons is equally shared between the atoms. It means that the electron density is symmetrically distributed over the two atoms.

Similar electronegativities

The red color represents electron density.

On the contrary, in the bond between dissimilar atoms such as HF, the electron density lies more towards the fluorine atoms as compared to the hydrogen atom. This tendency to attract a shared electron pair towards itself is termed electronegativity.

Very different electronegativities

The red color represents electron density.

Electronegativity cannot be directly measured and is calculated from other properties such as bond energies, ionization energies, and electron affinities of the bonded atoms. However, it is to be noted that electronegativity is a relative property of atoms in a molecule whereas properties such as ionization energies and electron affinities are properties of isolated atoms. 

Pauling scale is the most common scale for measuring electronegativity. The electronegativity values of various elements were calculated by Pauling, an American chemist (1901-1994), from the difference between the expected bond energies and their experimentally determined values.

Based on his calculations, he devised a scale for measuring electronegativity on which Fluorine is given an arbitrary standard value i.e. 4.0. It implies that fluorine is the most electronegative element of the periodic table. The electronegativity values for other elements were compared with that of fluorine. 

Factors Affecting Electronegativity

Electronegativity depends upon the atomic structure and the nature of atoms forming the bond. In general, small atoms are more electronegative as the bonded pair of electrons is closer to the nucleus and a greater force of electrostatic attraction follows. 

Similarly, the atoms with their valence shells nearly filled (that can form bonds too i.e., halogens) possess higher electronegativity than atoms with half or scarcely filled shells (i.e., alkali metals). It is the atomic number that explains this occurrence. In fact, a higher atomic number corresponds to stronger electrostatic attraction from the nucleus towards the valence electrons (and smaller atomic size).

The electronegativity of an element refers to its value in its most common oxidation state. The higher the oxidation state of an element, the more strongly will the electron pair be attracted. 

The chemical effects of increase in electronegativity with an increasing oxidation state can be noticed in the structures of oxides and halides, and in the acidity of oxides and oxoacids. CrO3 and Mn2O7 are acidic oxides having low melting points, whereas Cr2O3 is amphoteric and Mn2O3 is a basic oxide (with the decreasing oxidation state of metals).

The electronegativity of an atom also depends on the hybridization of the orbitals used in bonding. Electrons of s-orbitals are held more tightly than the electrons of p-orbitals. It implies that an orbital with more s-character will have a higher electronegativity value. 

The electronegativities for different hybridization schemes of an element follow the order:

χ (sp3) < χ (sp2) < χ (sp)

Trends of electronegativity

In the periodic table, electronegativity increases left to right across a period with the decreasing atomic size. Down a group, electronegativity decreases along with the increase in atomic size, with a few exceptions in the transition elements. The successive appearance of electronic shells at each step accounts for an increased shielding effect, causing the nuclear attraction to be reduced.

The difference in electronegativities of two elements can be linked to other properties such as dipole moment and bond energy. It is the difference in electronegativity between two bonded atoms that determines whether the covalent bond will be non-polar, polar, or even ionic in character.

If the electronegativity difference between the two bonded atoms is 0.4 or less, the bond is essentially non-polar. Contrary to this, if the electronegativity difference is between 0.4 to 1.7, the covalent bond will be polar in nature. In case the electronegativity difference is greater than 1.7, the bond will be termed ionic bond.

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So the bond between similar atoms has no electronegativity difference and is non-polar. Comparatively, a bond between different atoms is mostly polar. A difference of 1.7 units shows roughly equal contributions of the covalent and ionic bonds.

Exceptions in the Trends of Electronegativity

In general, electronegativity increases across a period and decreases down a group. It makes fluorine the most electronegative element, and conversely cesium and francium the least electronegative ones. However, here are a few exceptions to this general rule. 

Because of d-block contraction, gallium and germanium have higher electronegativities than aluminum and silicon respectively. 

As the 3d electrons are not very effective at shielding the increased nuclear charge, elements of the fourth period after the first transition metal series have unusually small atomic radii, which correlates to higher electronegativity values.

The unusually high electronegativity of lead is due to the increase in electronegativity with the oxidation state (for lead (+2) state goes better with the trend than (+4) state).

Scales of Electronegativity

There exist several scales to measure the electronegativities of elements:

Pauling Scale

Pauling first introduced the concept of electronegativity. He explained the electronegativity difference between two atoms (A – B) in terms of deviation (Δ) of the (A – B) bond energy from the geometric mean of (A – A) and (B – B) bond energies.

Electronegativity scale according to Pauling

It implies that if (A – B) bond energy is equal to the geometric mean of (A – A) and (B – B) bond energies, the two atoms will have the same electronegativity. That is because the two electrons are equally shared in a purely covalent bond, in this case. Hence, the (A – B) bond energy can be expressed as: 

EA-B = [EA-A x EB-B]½

However, Pauling noted that most of the (A – B) bond energies exceed the geometric mean as the atoms have different electronegativity values in general; that there exists some ionic character in the covalent bond.

Δ = EA-B – [EA-A x EB-B]½


  • Δ is the difference between the (A – B) bond energy and that of the geometric mean of (A – A) and (B – B).

The remaining excess energy is called ionic resonance energy. It can be used to determine the electronegativity difference between two atoms. If χA and χB represent the electronegativities of A and B, respectively, then the difference χA – χB can be related to Δ as:

χA – χB = (Δ / 96.5 kJ mol-1)½


  • The factor 96.5 kJ mol-1 converts Δ from kJ mol-1 to eV per molecule.

Electronegativity Formula and its Application

An example is the electronegativity difference of the (H-Cl) bond. The bond energies of (H – H), (Cl – Cl), and (H – Cl) are 436kJmol-1, 243kJmol-1, and 431kJmol-1 respectively.

To find the electronegativity difference: 

ΔHCl = EH-Cl – [EH-H x ECl-Cl]½

= 431 – [436 x 243]½


χH – χCl = (ΔHCl / 96.5 kJ mol-1)½

= [106 / 96.5]½

= 1.05 

The values devised by Pauling are given in the Pauling scale, which is the most common for measuring electronegativities. On this scale, fluorine has the highest electronegativity of 4.0.

Read more about the Pauling Scale of Electronegativity here.

Mulliken’s Scale

Electronegativity, in this scale, is partly determined by the tendency of an atom to gain an additional electron and partly by its tendency to retain those it already has. Hence it is taken as the average value of the first ionization energy and the first electron affinity. 

Both of the quantities are given a positive sign if the gain of electron involves release of energy and the loss of electron involves absorption of energy.


χ = (Ei + Eea) / 2

A complete electronegativity scale based on Mulliken’s approach cannot be constructed because the electron affinity is known for a few atoms only.

The values in terms of Pauling scale can be obtained from the equation: 

E.N = I.E + E.A (kJmol-1) / 544

Other Electronegativity Scales:

Allerd Rochow Scale

Another method of measuring electronegativities is given by Allred and Rochow. Its advantage lies in the fact that it can be applied to a greater number of elements. The reasoning followed here is that an atom attracts the electron density of a bond according to Coulomb’s law:

Force of attraction = Ze2 / 4πϵor2


  • Z is the effective nuclear charge
  • e is the charge on an electron
  • r is the radius of an atom

With the electronegativities empirically adjusted, the equation above gives:

χ = 0.359 (Z / r2) + 0.744

The numerical constants included, bring the values closer to those in the Pauling scale. This scale also follows that the atoms with smaller radii and higher effective nuclear charges are the most electronegative ones.

Allen Scale

Leland C. Allen developed a more recent scale for calculating “spectroscopic electronegativities”:

 χspec = mϵp + nϵs / (m + n)


  • m and n are the numbers of p and s electrons respectively
  • ϵp and ϵs are the corresponding average one-electron ionization enthalpies of an atom

The values of ϵp and ϵs are determined by high-resolution spectroscopy data. By this scale, Neon has the highest electronegativity, followed by fluorine, helium, and then oxygen.

Orbital Electronegativity

Electronegativity also depends on the nature of orbitals involved in bonding. As an example, an s-orbital lies closer to the nucleus than a p-orbital and so contributes to a more electronegative character. Therefore, the orbital electronegativities depend on the percentage of s and p characters present in the specified hybridization.

So for the carbon atom, the Pauling electronegativities are 2.48, 2.75, and 3.29 for sp3, sp2, and sp hybridizations respectively. The commonly used hybridization value for carbon is 2.5, based on its sp3 hybridization (tetrahedral).

Group Electronegativity

The electronegativity of an atom adjusted for substituents is known as group electronegativity. Thus, the group electronegativities of CH3 or CF3 groups will be the electronegativity of carbon adjusted in the presence of three H or F atoms respectively.

Electronegativity in groups

Several methods make use of atomic electronegativities, kinetic information, and other physical measurements to calculate the group electronegativities.


Electropositivity is the opposite of electronegativity. It is a measure of the metallic character of an element. Electropositivity can be understood as the ease of losing an electron. A higher electropositivity value indicates a stronger tendency to lose electrons.

The larger the size of an atom, the farther the valence electrons will be from the nucleus, and easier the removal will be. Contrary to electronegativity, electropositivity decreases from left to right across a period and increases down a group. 

It implies that the most electropositive atoms will be found at the bottom left of the periodic table, in contrast to the most electronegative elements on the top right of the periodic table.

Generally, the electropositive elements are more basic in nature. So the basic properties increase down a group. However, this trend does not hold for transition elements where the basicity decreases down a group. Electropositive elements possess the following properties:

I. React with water to form basic oxides and hydroxides.

II. More electropositive metals have a greater tendency to attract opposite charges; thus are not easily hydrated.

III. Have a small tendency to hydrolyze and form oxy-salts.

IV. Have little tendency to form complexes.

Applications of Electronegativity

The immediate application of electronegativities lies in the bond polarity. The greater the difference in electronegativity between two atoms, more polar the bond formed will be. The negative end of the dipole will be the atom with higher electronegativity. 

Several correlations are identified between infrared stretching frequencies of specific bonds and the electronegativities of the atoms involved. It follows that stretching frequencies depend partly on the strength of the bond.

There are other correlations between electronegativity and chemical shifts in NMR spectroscopy or isomer shifts in Mössbauer spectroscopy. Both the measurements depend on the s-electron density at the nucleus, hence proving electronegativity as “the ability of an atom in a molecule to attract electrons to itself”.

The greater the electronegativity difference between two bonded atoms, the more the electron density will be pulled toward the more electronegative atom. The lesser electronegative atom will therefore have a decreased electron density, and the protons attached to that atom will be de-shielded. As the E.N of the substituent increases, so does the extent of deshielding and the chemical shift. 

Concepts Berg

What if two atoms of equal electronegativity bond together?

If two atoms having the same electronegativities are bonded together, they will both pull the shared electron pair with equal force. The molecule will be non-polar with zero dipole moment.

Why does electronegativity increase across a period?

Across the period, the atomic number increases, and the atomic radius decreases. The nucleus is able to attract a shared pair of electrons with a stronger force.

What is electronegativity and how does it work?

Electronegativity is a measure of the tendency of an atom in a molecule to attract a bonded pair of electrons towards itself. Since the nucleus is positively charged and the electrons negatively charged, there is an electrostatic force of attraction between the nucleus and the pair of electrons.

What factors affect electronegativity?

The factors are atomic size, number of protons, oxidation state, and the type of hybridization involved in bonding.

Is H2O polar or nonpolar?

Water, H2O, is polar because there is a considerable electronegativity difference between the oxygen and hydrogen atoms. Moreover, due to the presence of two lone pairs resulting in a bent geometry, the dipole moment is not zero.

What is electronegativity and why is it important?

It is defined as the tendency of an atom to attract a shared pair of electrons. Its importance lies mostly in the determination of the polarity of a substance. HCl is polar because of an appreciable electronegativity difference between hydrogen and chlorine atoms. Organic substances are generally non-polar due to a lesser electronegativity difference between carbon and hydrogen atoms.

What are the applications of electronegativity?

The immediate application of electronegativity is the presence of dipole moment and bond polarity. It also affects the infrared stretching frequencies in infrared spectroscopy and the chemical shifts in NMR.

What is the electronegativity of CH4?

The electronegativity of C is 2.5 and that of H is 2.2 by the Pauling scale. The difference is roughly 0.3 which is considered non-polar for a bond.

How to determine electronegativity without the chart?

Electronegativities can be roughly predicted keeping in mind the factors that affect it, and its trends in the periodic table. Across a period (left to right), it increases with the increase in effective nuclear charge. Down a group, it decreases with increasing atomic radius.

What is the relationship between dipole moment and electronegativity?

It is the difference in electronegativity between two bonded atoms that determines the strength of charges present on the individual atoms. The dipole moment depends on the magnitude of this charge as well as the distance separating it. A higher difference corresponds to a larger dipole moment.