2.2: Electronegativity

Electronegativity is the measurement of an atom to compete for electrons in a bond. The higher the electronegativity, the greater its ability to gain electrons in a bond. Electronegativity will be important when we later determine polar and nonpolar molecules. Electronegativity is related with ionization energy and electron affinity. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted by the positive nucleus on the negative electrons. Therefore the electronegativity increases from bottom to top and from left to right.

Definitions of Electronegativity

Linus Pauling introduced the first electronegativity scale in 1932 in order to explain the extra stability of molecules with polar bonds.^[1] The electronegativity of an atom, represented by the Greek letter \(χ\) (chi), can be defined as the tendency of an atom to draw electrons to itself in a chemical bond. On the Pauling scale, the electronegativity difference between two atoms A and B was defined in terms of the dissociation energies E_d of the A-A, B-B, and A-B bonds:

Equation \(\PageIndex{1}\):

\[\chi_{A} - \chi_{B} = \sqrt{E_{d}(AB) - [E_{d}(AA) + E_{d}(BB)]/2} \nonumber \]

where the energies are expressed in electron volts. Pauling's scale of electronegativity ranges from Fluorine (most electronegative = 4.0) to Francium (least electronegative = 0.7). ^[2,3] The polarity of bonds and assignment of formal charges is predicted by electronegativity differences.

Figure \(\PageIndex{1}\): Table of Pauling electronegativities from [Wikipedia]

While directly relevant to the strength of chemical bonds, the Pauling definition of electronegativity has a significant limitation: calculating each value requires data from several compounds in which specific atoms are bonded. This means that the scale cannot be successfully applied to all situations. To overcome this limitation, alternative electronegativity scales were developed based on different thermochemical measurements or calculations. These other scales are listed in Table \(\PageIndes{1}\) below. More thorough descriptions of how values in each scale are calculated are described on Wikipedia's Electronegativity page and a concise history of electronegativity scales is summarized in ref 4.

Table \(\PageIndex{1}\): Electronegativity Scales and Their Defining Characterisitics

Scale Developer	Year first described	Characteristics
Linus Pauling^[1]	1932	Based on bond energies
Robert S. Mulliken^[5,6]	1934	Based on valence electron properties of atoms (electron affinity and ionization energy)
Louis Allred & Eugene G. Rochow^[7]	1958	Based on electrostatic force (effective nuclear charge)
Robert T. Sanderson^[8]	1983	Based on atomic electron density
Ralph G. Pearson^[9]	1985	Related to Hard-Soft Acid-Based theory
Leland C. Allen, Joseph Mann, Terry L. Meek^[10,11,12]	1989, 2000	Average ionization energies of valence shell electrons, configuration energies (CE)\[\mathrm{CE}=\frac{n \varepsilon_{s}+m \varepsilon_{p}}{n+m} \nonumber \] \(n = \text{number of s electrons}\) \(m = \text{number of p electrons}\) \(\varepsilon_{s}, \varepsilon_{p} = \text{experimental 1-electron s and p energies}\)

The primary advantage to the scale developed by Mann, Meek, and Allen^[12] is that it is based on configuration energy (CE), the average ionization energies of valence electrons in ground state free atoms. A scale based on ionization energies can be calculated more directly for any element. However, a critique of all electronegativity scales, including this one, is that they are all based on assumptions that fail in some cases.

There is a relationship between electronegativity and atomic size because both are related to ionization energy. Like ionization energy, there is a general trend across the periodic table for both electronegativity and atomic size; in general, smaller atoms toward the top right-hand side of the periodic table have greater electronegativity values and smaller atomic size.

Electronegativity Trends

Electronegativity is a qualitative property, and although there is no standardized method for calculating electronegativity the most common scale for quantifying electronegativity is the Pauling scale (Table A2), named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due to the qualitative nature of electronegativity. Electronegativity values for each element can be found on certain periodic tables (Figure \(\PageIndex{2}\).

Figure \(\PageIndex{2}\): Periodic Table with Electronegativity Values

Electronegativity measures an atom's tendency to attract and form bonds with electrons. This property exists due to the electronic configuration of atoms. Most atoms follow the octet rule (having the valence, or outer, shell comprise of 8 electrons). Because elements on the left side of the periodic table have less than a half-full valence shell, the energy required to gain electrons is significantly higher compared with the energy required to lose electrons. As a result, the elements on the left side of the periodic table generally lose electrons when forming bonds. Conversely, elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a complete valence shell of 8 electrons. The nature of electronegativity is effectively described thus: the more inclined an atom is to gain electrons, the more likely that atom will pull electrons toward itself.

• From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one.
• From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.
• Important exceptions of the above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values.
• As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.

According to these two general trends, the most electronegative element is fluorine, with 3.98 Pauling units.

Figure \(\PageIndex{3}\): Periodic Table showing Electronegativity Trend (LibreText)

Acknowledgment

• Matthew Salem (UC Davis) (Pauling Electronegativity)
• Introduction to Inorganic Chemistry Wikibook (https://chem.libretexts.org/Bookshel..._Bond_Strength)
• Kathryn Haas (Duke) (https://chem.libretexts.org/@go/page/219071)