Electronegativity Definition and Trend

This entry was posted on by (updated on )
Electronegativity trend of the periodic table
Periodic table showing the trend from low to high of the element’s electronegativity.

Electronegativity is a measure of an atom’s ability to attract and hold onto electrons within a chemical bond. A high electronegativity value means an atom readily attracts electrons to form a chemical bond with another atom. A low electronegativity value means an atom readily donates electrons to form a bond or is electropositive.

While there are charts of electronegativity values for elements of the periodic table, there is no true single electronegativity value for an atom. Rather, it depends on the other atoms in a molecule and also depends on the nuclear charge and number of electrons. The most common method of calculating electronegativity is the Pauling scale, which was proposed by Linus Pauling. The Pauling scale runs from 0.79 to 3.98. The Pauling scale is dimensionless, but sometimes the values are cited in Pauling units.

Key Points: Electronegativity

  • Electronegativity is a measure of an atom’s ability to attract electrons in a chemical bond.
  • Fluorine has the highest electronegativity (4.0), while cesium has the lowest (0.7) on the Pauling scale.
  • Electronegativity increases across a period and decreases down a group.
  • Differences in electronegativity determine bond polarity and molecular behavior.
  • Atoms with comparable electronegativity form covalent bonds, while atoms with large electronegativity differences form ionic bonds,
  • Scales include Pauling, Mulliken, Allred-Rochow, and Allen.

Most Electronegative and Most Electropositive Elements

The most electronegative element is fluorine, with an electronegativity value of 3.98 on the Pauling scale. The least electronegative or most electropositive element is cesium, which has a value of 0.79. However, francium is probably even more electropositive than cesium because it has a higher ionization energy. Francium’s electronegativity value is estimated to be around 0.79, but it has not been measured empirically.

Electronegativity and Chemical Bonding

Comparing electronegativity values allows prediction of the type of chemical bond two atoms will form. Atoms with the same electronegativity values (e.g., H2, N2) form covalent bonds. Atoms with slightly different electronegativity values (e.g., CO, H2O) form polar covalent bonds. All hydrogen halides (e.g., HCl, HF) form polar covalent bonds. Atoms with very different electronegativity values (e.g., NaCl) form ionic bonds. Note that electronegativity does not help in predicting whether or not a chemical bond will actually form. Argon has a high electronegativity value, yet it’s a noble gas that forms few chemical bonds.

Electronegativity Periodic Table Trend

Electronegativity follows a trend (periodicity) on the periodic table. The trend is shown in the graphic (which is also available as a PDF for printing).

  • Electronegativity increases moving left to right across a period, from the alkali metals to the halogens. The noble gases are an exception to the trend.
  • Electronegativity decreases moving down a periodic table group. This is because the distance between the nucleus and the valence electrons increase.
  • Electronegativity follows the same general trend as ionization energy. Elements with low electronegativities tend to have low ionization energies. Similarly, an atom with a high electronegativity tends to have a high ionization energy.

Examples and Exceptions

Electronegativity trends provide powerful insights, but examples and exceptions clarify how elements behave.

Examples:

  • High Electronegativity: Fluorine (4.0) strongly attracts bonding electrons, making it the most electronegative element.
  • Low Electronegativity: Cesium (0.7) readily loses its single valence electron, forming ionic bonds with electronegative elements.
  • Intermediate Electronegativity: Carbon (2.5) balances electronegativity, forming nonpolar or polar covalent bonds depending on its partner.

Exceptions:

  • Hydrogen: Though a nonmetal, hydrogen’s electronegativity (2.1) rivals some metals, making it versatile in forming bonds.
  • Transition Metals: d-orbitals and the lanthanide contraction complicate electronegativity trends, leading to less predictable behavior.

Factors Influencing Electronegativity

Electronegativity trends depend on atomic structure and properties. Several factors combine to explain why certain elements attract electrons more strongly than others.

  1. Atomic Number:
    • As an atom’s nuclear charge increases, its ability to pull electrons also grows. This explains why electronegativity increases across a period.
  2. Atomic Radius:
    • Smaller atoms bring electrons closer to the nucleus, intensifying the attractive force. This makes elements like fluorine highly electronegative.
  3. Electron Shielding:
    • Inner electron shells block part of the nuclear charge from reaching outer electrons. As a result, larger atoms with more shielding, like cesium, have lower electronegativity.
  4. Oxidation State:
    • When an atom loses electrons (achieving a higher oxidation state), its nuclear pull increases, enhancing electronegativity.

These factors work together to shape electronegativity trends, which chemists rely on to predict bonding and molecular behavior.

Electronegativity Scales

Electronegativity reflects an atom’s ability to attract electrons, but scientists measure it differently depending on their approach. While the Pauling scale is the most widely used, other scales like the Mulliken, Allred-Rochow, and Allen scales offer unique perspectives. Each method highlights different aspects of an atom’s behavior.

  1. Pauling Scale:
    • Linus Pauling developed this scale by comparing bond dissociation energies. It is the most common because it closely matches experimental data for bond polarities and is straightforward for practical chemistry applications.
    • Values are dimensionless and range from 0.7 (cesium) to 4.0 (fluorine).
  2. Mulliken Scale:
    • This scale averages an atom’s ionization energy and electron affinity, expressing values in electron volts (eV). It provides an absolute measure of electronegativity.
  3. Allred-Rochow Scale:
    • This scale focuses on effective nuclear charge and how strongly the nucleus attracts valence electrons. It emphasizes electrostatic forces at play in atomic bonding.
  4. Allen Scale:
    • This method calculates electronegativity using the average energy of valence electrons. It is helpful for theoretical studies but less common in practical use.

Real-World Applications of Electronegativity

Electronegativity shapes how atoms interact, influencing chemistry at every level. Its importance spans multiple fields:

  1. Chemical Bonding:
    • Chemists use electronegativity to predict bond types. Large differences lead to ionic bonds, while small differences result in covalent bonds.
  2. Material Science:
    • Engineers design materials by manipulating atomic interactions, such as using semiconductors with specific electronegativity values.
  3. Medicinal Chemistry:
    • Drug developers analyze molecular polarity to optimize hydrogen bonding and molecular interactions in the body.
  4. Environmental Chemistry:
    • Electronegativity helps researchers study pollutant behavior, such as the high reactivity of halogens in atmospheric chemistry.
  5. Biological Systems:
    • Water’s polarity, driven by oxygen’s high electronegativity, underpins many biological processes, including protein folding and enzyme activity.

References

  • Jensen, William B. (1996). “Electronegativity from Avogadro to Pauling: Part 1: Origins of the Electronegativity Concept.” J. Chem. Educ. 73 (1): 11-20. ACS Publications. doi:10.1021/ed073p11
  • Jolly, William L. (1991). Modern Inorganic Chemistry (2nd ed.). New York: McGraw-Hill. ISBN 978-0-07-112651-9.
  • Mullay, J. (1987). Estimation of atomic and group electronegativities. Structure and Bonding. 66:1–25. doi:10.1007/BFb0029834. ISBN 978-3-540-17740-1.
  • Pauling, Linus (1932). “The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms.” J. Am. Chem. Soc. 54 (9): 3570-3582. doi:10.1021/ja01348a011
  • Pauling, Linus (1960). The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Mode (3rd ed.). Cornell University Press. ISBN 978-0-8014-0333-0.