Bond polarity and molecular polarity are related but distinct concepts, both influenced by electronegativity. Bond polarity refers to the distribution of electrical charge between two atoms in a bond. It arises from the difference in electronegativity between the atoms. A larger difference leads to greater bond polarity, with electrons being more attracted to the atom with higher electronegativity. In contrast, molecular polarity is determined not just by bond polarity but also by the shape and symmetry of the molecule. A molecule is polar if it has polar bonds and an asymmetrical structure, resulting in an uneven distribution of charge. For example, water (H₂O) is a polar molecule because it has polar O-H bonds and a bent shape, leading to an uneven charge distribution. Conversely, carbon dioxide (CO₂) is non-polar despite having polar C=O bonds, as its linear symmetrical shape allows the bond polarities to cancel out. Thus, while bond polarity is a direct result of electronegativity differences, molecular polarity also depends on the molecular geometry.

Electronegativity is central to understanding hydrogen bonding, a special type of dipole-dipole interaction. Hydrogen bonding occurs when a hydrogen atom bonded to a highly electronegative atom (like nitrogen, oxygen, or fluorine) is attracted to another electronegative atom in a nearby molecule. The high electronegativity of these atoms causes a significant polarity in their bonds with hydrogen, leading to the partial positive charge on hydrogen being strongly attracted to the partial negative charges on electronegative atoms in other molecules. This interaction is particularly significant in substances like water, where hydrogen bonding contributes to its unusually high boiling and melting points, high surface tension, and the expansion of water upon freezing. In biological molecules like DNA and proteins, hydrogen bonding is crucial for maintaining structure and function. The strength and directionality of hydrogen bonds, influenced by the electronegativity of the involved atoms, play a vital role in the three-dimensional structure and properties of these complex molecules. Therefore, electronegativity not only helps in predicting the existence of hydrogen bonds but also in understanding their profound effects on the physical and chemical properties of substances.

Electronegativity plays a subtle but important role in the interpretation of Lewis structures, which are diagrams that show the bonding between atoms and the lone pairs of electrons in molecules. While Lewis structures primarily focus on illustrating how atoms are bonded, electronegativity helps in understanding the nature of these bonds and the distribution of electron density. In molecules with atoms of different electronegativities, Lewis structures can be used to predict areas of higher electron density around the more electronegative atoms. This can lead to an understanding of the molecular polarity. For instance, in a water molecule (H₂O), the Lewis structure shows two lone pairs on the oxygen atom, and the higher electronegativity of oxygen compared to hydrogen indicates that the electron density is more concentrated around oxygen, leading to a polar molecule. Additionally, electronegativity differences can suggest the likelihood of multiple bonding and the presence of polar bonds, which are key considerations in predicting molecular shapes and reactivities.

Electronegativity can provide insights into the reactivity of elements, particularly in their ability to form chemical bonds. Elements with high electronegativity, such as fluorine and oxygen, are often more reactive because they have a strong tendency to attract electrons from other atoms. This makes them excellent oxidizing agents. In contrast, elements with low electronegativity, such as alkali metals, are more likely to donate electrons, making them highly reactive as reducing agents. However, it's important to note that while electronegativity is a useful guide, reactivity is also influenced by other factors like atomic size, ionization energy, and electron affinity. For example, noble gases have high electronegativities but are generally unreactive due to their stable electronic configurations. Therefore, while electronegativity is a helpful indicator, it should be considered alongside other atomic properties for a comprehensive understanding of an element’s reactivity.

Electronegativity significantly influences the physical properties of compounds, such as melting point, boiling point, and solubility. In ionic compounds, where electronegativity differences are high (e.g., sodium chloride), there is a strong electrostatic attraction between the ions, resulting in high melting and boiling points. Such compounds are often soluble in polar solvents like water but insoluble in non-polar solvents. Conversely, in covalent compounds, where electronegativity differences are lower, the intermolecular forces are weaker, typically resulting in lower melting and boiling points. Nonpolar covalent compounds (e.g., methane) are generally insoluble in water but soluble in non-polar solvents. Polar covalent compounds (e.g., water) can have varying solubilities depending on the polarity and hydrogen bonding, often exhibiting higher solubilities in polar solvents. Thus, the difference in electronegativity between atoms in a molecule directly impacts the strength of the bonds and, consequently, the physical properties of the compound.