By Riti Gupta Updated Aug 30, 2022
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Electronegativity is the driving force that determines how strongly an atom attracts electrons. The greater an atom’s electronegativity, the more it pulls shared electrons toward itself. This principle underpins the variety of chemical bonds we encounter.
When one atom is markedly more electronegative than its partner, two scenarios arise. It may fully capture an electron from the other atom, forming an ionic bond. Alternatively, it can simply draw the shared electrons closer, creating a polar covalent bond. Bonds involving highly electronegative atoms such as oxygen or fluorine typically exhibit this polar character, with the electronegative atom bearing a partial negative charge (δ‑) and its partner a partial positive charge (δ+).
Bonds fall along a spectrum from purely nonpolar to fully polar. An ionic bond represents the extreme, where one atom’s electronegativity difference is so large that it takes an electron entirely. A nonpolar covalent bond occurs when the two atoms share electrons equally, as seen in homonuclear diatomic molecules like H₂ or Cl₂.
For bonds that lie between these extremes, the difference in electronegativity (ΔEN) dictates the bond type. The following table summarizes the commonly accepted thresholds:
| Bond Type | Electronegativity Difference (ΔEN) |
|---|---|
| Pure Covalent | < 0.4 |
| Polar Covalent | 0.4 – 1.8 |
| Ionic | > 1.8 |
Thus, the distinction between polar and nonpolar bonds hinges on ΔEN.
A molecule may contain polar covalent bonds yet remain nonpolar if the bond dipoles cancel each other out due to molecular symmetry. Water (H₂O) is a classic polar compound: its bent geometry leaves a net dipole, enabling extensive hydrogen bonding. In contrast, boron trifluoride (BF₃) features three polar B–F bonds arranged in a planar trigonal shape that cancel out, rendering the molecule overall nonpolar.
Polarity profoundly influences how molecules interact, especially in solution. For instance, sugars dissolve readily in water because both molecules possess polar functional groups that form hydrogen bonds. Water’s oxygen atom carries a partial negative charge, while the hydrogen atoms bear partial positive charges, creating an environment that favors hydrogen bonding with sugar’s hydroxyl (–OH) groups.
Conversely, oils consist largely of nonpolar C–H bonds. The small electronegativity difference between carbon and hydrogen results in negligible dipole moments, preventing hydrogen bonding. Thus, oil molecules do not interact favorably with water’s polar network, leading to immiscibility.
By examining a compound’s structural features and the nature of its bonds, chemists can predict whether it will exhibit polarity and, consequently, its behavior in various chemical contexts.
