Electronegativity: Trends across the Periodic Table

Definition and Significance of Electronegativity

Electronegativity is a fundamental concept in chemistry that measures an atom's ability to attract electrons towards itself in a chemical bond. It plays a crucial role in determining the type of bond formed between atoms, whether it is ionic, covalent, or polar covalent. The electronegativity value of an element is influenced by factors such as atomic size, electron affinity, and the number of valence electrons.

Increasing Electronegativity from Left to Right

Across a period in the periodic table, electronegativity values generally increase from left to right. This trend can be attributed to the increasing effective nuclear charge experienced by the electrons in the outermost shell. As the number of protons in the nucleus increases, the attractive force on the electrons becomes stronger, resulting in a higher electronegativity value.

Decreasing Electronegativity from Top to Bottom

Within a group on the periodic table, electronegativity values tend to decrease from top to bottom. This trend is primarily influenced by the increasing distance between the nucleus and the outermost shell as you move down the group. The larger atomic size and shielding effect from inner electron shells reduce the attractive force on the outermost electrons, leading to a lower electronegativity value.

Example: Most Electronegative and Least Electronegative Elements

Fluorine holds the title for the most electronegative element on the periodic table. With its small atomic size and high effective nuclear charge, fluorine has a strong ability to attract electrons towards itself in a chemical bond. On the other end of the spectrum, francium possesses the lowest electronegativity value. As the heaviest alkali metal, francium has a larger atomic size and weaker attractive force on the outermost electrons, resulting in a lower electronegativity value.